IABOMTORY^PERIMENTS 


IN 


GENERAL  CHEMISTRY 


H-B' NORTH 


D.^&N  NOSTRAND  COMPANY 


Edmund  O'Neill 


Laboratory  Experiments 


in 


General  Chemistry 


BY 
H.    B.    NORTH,    PH.G.,    D.Sc. 

Associate  Professor  of  Chemistry  in  Rutgers  College 

Member  of  the  American   Chemical  Society 

A  merican  Electrochemical  Society 

Societe  chimique  de  France 


36  ILLUSTRATIONS 


NEW    YORK 

D.    VAN    NOSTRAND    COMPANY 

25  PARK  PLACE 

1913 


COPYRIGHT,  1913, 

BY 

D.  VAN  NOSTRAND  COMPANY. 
IN 


Stanbq>c 

F.    H.GILSON   COMPANY 
BOSTON,  U.S.A. 


PREFACE. 


THIS  manual  is  designed  to  cover  a  laboratory  course 
in  General  Chemistry  given  in  connection  with  a  series 
of  experimental  lectures.  It  contains  five  hundred  care- 
fully chosen  experiments  on  the  more  common  elements 
and  is  so  arranged  that  it  can  be  used  in  connection 
with  any  good  text-book.  The  work  includes  a  large 
number  of  experiments  similar  to  those  found  in  other 
manuals  and,  in  addition,  numerous  more  advanced 
experiments  which,  to  the  author's  knowledge,  have 
never  before  appeared  in  a  laboratory  manual  in 
General  Chemistry. 

It  is  not  supposed  that  any  one  student  will  perform 
all  of  these  experiments.  The  reason  for  the  large 
number  is  rather  that  experiments  may  be  chosen  to 
meet  the  needs  of  the  various  classes  of  students.  In 
the  author's  laboratory  an  assignment  of  experiments 
for  each  laboratory  period  is  posted  on  the  bulletin 
board.  A  number  of  the  simpler  experiments  are 
selected  for  the  beginners  while  the  more  advanced 
and  consequently  more  difficult  exercises  are  assigned 
to  those  who  have  had  previous  chemical  training. 
In  order  to  better  facilitate  this  method  of  assignment, 
all  experiments  have  been  numbered  consecutively. 

In  writing  this  book,  the  author  has  attempted  to 
word  each  and  every  experiment  in  such  a  way  as  to 


889773 


v  PREFACE 

make  it  impossible  for  the  student  to  mistake  the 
exact  meaning.  A  preliminary  edition  of  the  work  has 
been  in  use  in  the  author's  laboratory  for  one  year  and 
has  proved  most  satisfactory. 

H.  B.  N. 

NEW  BRUNSWICK,  N.  J., 
Aug.  i,  1913. 


CONTENTS. 


CHAPTER  PAGE 

I.   CAUSES  or  CHEMICAL  CHANGE i 

II.  HYDROGEN 3 

III.  OXYGEN  AND  OZONE 15 

IV.  WATER  AND  HYDROGEN  PEROXIDE 23 

V.  THE  HALOGENS 35 

VI.  ACIDS,  BASES  AND  SALTS 53 

VII.  NITROGEN 58 

VIII.  OXIDATION  AND  REDUCTION 73 

IX.   SULPHUR 77 

X.   CARBON 87 

XI.  SILICON  AND  BORON 98 

XII.   PHOSPHORUS,  ARSENIC,  ANTIMONY  AND  BISMUTH 104 

XIII.  THE  ALKALIES  AND  AMMONIUM 118 

XIV.  THE  ALKALINE  EARTHS 128 

XV.   MAGNESIUM,  ZINC,  CADMIUM  AND  MERCURY 137 

XVI.   COPPER,  SILVER  AND  GOLD 144 

XVII.  TIN  AND  LEAD 152 

XVIII.  ALUMINUM  AND  CHROMIUM 157 

XIX.   MANGANESE 166 

XX.  IRON,  COBALT  AND  NICKEL 171 

XXI.  PLATINUM 179 

APPENDIX 181 

Correction  of  Gas  Volumes 181 

Chemical  Arithmetic 186 

Tables 193 


LABORATORY   EXPERIMENTS 
IN  GENERAL  CHEMISTRY 


CHAPTER  I. 

CHEMICAL   CHANGES   AND   THE   AGENCIES 
WHICH  PRODUCE  THEM. 

1.  Carefully   weigh    a    small   porcelain    evaporating 
dish;    then  weigh  into  ;  it  £?g,ctjy  5  ,gms.  of  powdered 
iron.    Place  the  dish  Qii  £•  ring,  stand  *uid  heat  strongly 
for  about  10  minutes.    ,A}low:  tc  G(?oJ;,  ^heri'  re- weigh. 

Compare  the  weights  of 'tile  itan.  Has  the  iron  in- 
creased or  decreased  in  weight?  How  much?  What  is 
the  reason  for  this  change?  Through  what  agency  has 
this  change  been  brought  about  ? 

2.  To  a  few  cubic  centimeters  of  a  solution  of  sodium 
chloride  (NaCl)  in  a  test  tube  add  a  little  silver  nitrate 
(AgN03)  solution.     Observe  that  the  compound  formed 
is  pure  white  in  color. 

Allow  the  test  tube  containing  the  white  precipitate 
to  stand  on  the  desk  for  an  hour,  or  until  the  end  of  the 
laboratory  period.  If  convenient,  place  it  in  the  direct 
rays  of  the  sun.  Has  any  change  in  appearance  taken 
place?  If  so,  what  has  caused  this  change? 

3.  Introduce  a  little  powdered  sodium  bicarbonate 
(NaHC03)  into  a  test  tube  and  treat  with  water.    In 
another    tube    treat    some    potassium    acid    tartrate 


2  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

(KHC4H406)  with  water.  Have  any  changes  taken 
place?  If  so,  are  they  physical  or  chemical  changes? 

In  a  dry  porcelain  mortar  rub  together  about  equal 
portions  of  sodium  bicarbonate  and  potassium  acid 
tartrate.  Do  you  notice  any  change?  Now  place  the 
mixture  in  a  dry  beaker  and  add  water.  What  phenom- 
enon do  you  observe  ?  What  agency  is  the  cause  of  this 
change? 

Summary.  What  is  the  difference  between  a  physical 
and  a  chemical  change?  In  what  does  the  study  of 
chemistry  differ  from  that  of  physics  ?  Mention  several 
physical  changes  and  several  chemical  changes. 

What  three  agencies  of  chemical  change  have  been 
studied  in  the  preceding  eipedinen  ts  ?  Mention  another 
important  agency  of  chemical  change. 

Mention  some  :of  the  phenomena  which  indicate 
chemical  change. 


CHAPTER  II. 
HYDROGEN  (H;  i). 

4.  Preparation.    Place  a  small  piece  of  zinc  (Zn)  in  a 
test  tube  and  add  a  few  cubic  centimeters  of  dilute 
hydrochloric  acid  (HC1).    What  happens?     Bring  the 
mouth  of  the  test  tube  to  a  flame.    What  happens  ? 

In  like  manner  try  the  action  of  hydrochloric  acid  on 
small  pieces  of  magnesium  (Mg),  iron  (Fe)  and  alu- 
minum (Al).  Try  the  action  of  dilute  sulphuric  acid 
(H2S04)  on  zinc,  magnesium,  iron  and  aluminum.  Heat 
gently  if  necessary. 

Can  you  make  a  general  statement  covering  all  these 
cases  ? 

5.  Carefully  dry  a  small  piece  of  metallic  sodium  (Na) 
and  wrap  it  in  a  piece  of  dry  filter  paper.-    Fill  a  test 
tube  with  water  and  invert  it  in  a  dish  of  water,  holding 
the  mouth  of  the  tube  under  the  surface.     By  means  of 
pincers  quickly  introduce  the  piece  of  sodium  wrapped 
in  paper  into  the  mouth  of  the  test  tube  under  the  water. 
What  do  you  observe?    Test  the  gas  formed. 

6.  Place  a  few  small  pieces  of  metallic  aluminum  in  a 
test  tube  and  add  a  few  cubic  centimeters  of  a  strong 
solution  of  potassium  hydroxide  (KOH)  or  sodium  hy- 
droxide (NaOH).    Heat  in  the  Bunsen  flame.     Explain 
the  action. 

Repeat  the  experiment,  using  metallic  zinc  instead  of 
aluminum. 

3 


EXPERIMENTS   IN   GENERAL   CHEMISTRY 


7.  Arrange  an  apparatus  as  shown  in  Fig.  i.  The 
flask  contains  100  cc.  of  water  and  a  few  pieces  of  broken 
tile  or  pumice  to  prevent  "  bumping."  The  iron  tube 
should  be  partially  filled  with  iron  filings.  Make  all 
connections  tight. 

Heat  the  iron  pipe  strongly  and  at  the  same  time  boil 
the  water  in  the  flask.  When  the  air  has  been  driven 
from  the  apparatus,  place  an  inverted  test  tube  filled 
with  water  over  the  end  of  the  delivery  tube.  What  gas 


FIG.  i. 

collects  in  the  test  tube  ?    Test  it  by  bringing  to  a  flame. 
How  have  the  iron  filings  changed  in  appearance? 

8.  Properties.  Arrange  a  hydrogen  generator  as 
shown  in  Fig.  2,  and  at  the  same  time  prepare  the 
tubes  necessary  in  Exps.  n  and  12.  (The  apparatus 
must  be  submitted  to  the  approval  of  the  instructor 
before  going  on  with  the  experiment.)  Place  25  gms. 
of  zinc  (Zn)  in  the  flask;  add  enough  water  to  barely 
cover  the  metal.  Now  add  concentrated  hydrochloric 
acid  (HC1)  through  the  thistle  tube,  a  little  at  a  time, 
until  a  brisk  effervescence  takes  place.  (CAUTION! 
Never  bring  a  flame  near  a  hydrogen  generator.) 


HYDROGEN  5 

To  test  the  gas,  collect  a  test  tube  full  over  water  and 
then  quickly  bring  the  mouth  of  the  test  tube  to  a  flame. 
If  an  explosion  follows,  the  hydrogen  is  impure.  If  the 
hydrogen  is  pure,  it  will  burn  quietly.  Test  the  gas  at 
intervals  until  it  is  found  to  be  pure.  Before  making  a 
test,  be  sure  that  there  is  no  hydrogen  burning  in  the 
test  tube. 


FIG.  2. 

When  the  gas  has  been  ascertained  to  be  pure,  collect 
several  bottles  over  water  in  the  manner  shown  in  the 
drawing.  These  bottles  of  gas  are  to  be  used  in  the 
following  experiments. 

9.  Why  is  hydrogen  always  kept  in  inverted  bottles? 
Pour  hydrogen  upward  into  an  empty  bottle.  Then 
test  both  bottles  to  see  if  they  contain  any  of  the 
gas. 

What  is  meant  by  the  term  "  vapor  density  "  ?  What 
is  the  standard  of  vapor  density?  What  is  the  vapor 
density  of  air? 


6  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

10.  Thrust  a  burning  splinter  upward  into  an  inverted 
bottle  of  hydrogen.     Is  the  flame  extinguished?    Does 
the  hydrogen  burn  at  the  mouth  of  the  bottle?     Care- 
fully withdraw  the  splinter  so  that  it  will  ignite  again. 
Repeat  several  times. 

Does  hydrogen  support  combustion?  What  do  you 
understand  by  the  term  "  combustion  "? 

11.  Remove  the  delivery   tube  from   the  hydrogen 
generator  and  in  its  place  insert  a  short  piece  of  tubing 
drawn  to  a  point  at  the  outer  end.    Add  a  little  more 
HC1  to  the  generator  if  necessary.     Wrap  a  towel  about 
the  generator  and  then  ignite  the  gas  issuing  from  the 
glass  tip.     Notice  the  color  of  the  flame  immediately. 
(After  burning  some  time,  the  flame  will  become  yellow, 
due  to  the  sodium  in  the  glass.)     Invert  a  clean  dry 
beaker  for  a  moment  over  the  burning  jet  of  hydrogen. 
What  collects  on  the  inside  of  the  beaker  ?     What  is  the 
product  of  combustion  of  hydrogen? 

12.  Attach  a  wash  bottle  and  delivery  tube  to  the 
generator 'as  shown  in  Fig.  3.    Add  2  or  3  cc.  of  potas- 
sium permanganate  (KMnC^)  solution  to  a  beaker  half 
full  of  water.     By  means  of  the  delivery  tube,  pass 
hydrogen  through   the   solution.     Do  you  notice  any 
change  ? 

Remove  the  hydrogen  generator.  Then  generate 
hydrogen  in  the  beaker  containing  the  solution  by 
dropping  in  a  few  pieces  of  zinc  (Zn)  and  then  adding 
a  few  cubic  centimeters  of  concentrated  sulphuric  acid 
(H2S04).  If  no  gas  is  generated,  add  more  acid.  Allow 
to  stand  for  several  minutes.  What  change  takes  place  ? 
Why  does  hydrogen  cause  a  change  in  this  case  and  not 
in  the  previous  one? 


HYDROGEN 


Repeat  using  a  solution  of  potassium  dichromate 
(K2Cr2O7)  instead  of  potassium  permanganate.  Describe 
the  results. 

o 


FIG.  3. 

13.  Other  Product  of  the  Action  of  an  Acid  on  a 
Metal.  Filter  the  solution  formed  in  the  hydrogen 
generator  and  carefully  evaporate  the  clear  solution  to 
dryness  in  a  porcelain  evaporating  dish.  Describe 
the  nature  of  the  material  left  in  the  dish.  (Hood.) 

Summary.  What  four  general  methods  for  the  pro- 
duction of  hydrogen  have  been  studied?  Mention  one 
other  good  method  for  the  preparation  of  this  gas. 

For  what  is  hydrogen  taken  as  the  standard  and  why 
is  it  so  taken?  What  is  meant  by  molecular  hydrogen 
and  by  nascent  hydrogen?  Does  the  latter  differ  from 
the  former  in  chemical  properties?  If  so,  which  is  the 
more  active? 

Problems.*  (a)  To  prepare  30  gms.  of  hydrogen  by  the  action 
of  HC1  on  zinc,  what  weight  of  zinc  is  necessary?  What  is  the 
volume  of  the  30  gms.  of  hydrogen? 

*  See  "  Chemical  Arithmetic,"  Appendix,  page  186. 


8 


EXPERIMENTS   IN   GENERAL   CHEMISTRY 


(b)  How  many  liters  of  hydrogen  can  be  produced  by  the 
action  of  an  excess  of  H2SO4  on  215  gms.  of  metallic  zinc? 

(c)  If  equal  weights  of  zinc,  aluminum,  iron  and  magnesium 
are  dissolved  in  HC1,  which  will  produce  the  greatest  amount 
of  hydrogen? 

REDUCTION  BY  HYDROGEN. 

(Quantitative.) 

14.  Arrange  an  apparatus  consisting  of  a  hydrogen 
generator,  wash  bottle,  calcium  chloride  tube  (a),  hard 
glass  tube  (b)  and  a  second  calcium  chloride  tube  (c),  as 
shown  in  Fig.  4.  Both  of  the  calcium  chloride  tubes 

o 


FIG.  4. 


should  be  filled  with  dry  granulated  calcium  chloride, 
CaCl2.  Prepare  caps  for  the  ends  of  c  as  shown  in  the 
drawing  so  that  this  tube  may  be  weighed  without 
danger  of  its  contents  absorbing  moisture  from  the  air. 
The  caps  can  be  conveniently  made  of  i-inch  pieces 
of  small  rubber  tubing,  one  end  being  plugged  with  a 
short  piece  of  glass  rod. 

Thoroughly  clean  the  hard  glass  tube  and  then  weigh 

it  accurately.     Now  introduce  about  5  gms.  of  ferric 

oxide  (Fe203)*  into  the  tube  as  near  the  center  as  pos- 

*  Cupric  oxide,  CuO,  may  be  substituted  for 


HYDROGEN  9 

sible,  and  again  weigh.  The  difference  in  the  two 
weighings  represents  the  amount  of  Fe20s  taken. 

Cap  the  ends  of  c  and  weigh  accurately.  Then  con- 
nect the  apparatus  as  shown  in  the  drawing,  first,  how- 
ever, introducing  about  30  gms.  of  zinc  into  the  flask. 
Through  the  thistle  tube  add  enough  water  to  cover  the 
zinc  and  then  sufficient  concentrated  HC1  to  produce 
a  brisk  evolution  of  hydrogen.  Wrap  a  towel  about 
the  generator  and  then  test  the  gas  being  evolved  from 
the  apparatus.  To  do  this,  hold  an  inverted  test  tube 
over  the  end  of  c  from  which  gas  is  issuing,  thus  collect- 
ing the  test  tube  full  of  hydrogen  by  displacement  of  air; 
then  quickly  bring  the  mouth  of  the  test  tube  to  a  flame. 
If  the  gas  is  pure  it  will  burn  quietly.  When  the  hy- 
drogen coming  from  the  exit  tube  is  pure,  apply  heat  to 
the  middle  of  the  hard  glass  tube  by  means  of  a  Bunsen 
burner. 

The  hydrogen,  coming  from  the  generator,  is  freed 
from  HC1  fumes  by  the  wash  bottle  of  water,  and  is  then 
dried  by  the  calcium  chloride  tube  a.  The  Fe203  heated 
in  the  dry  hydrogen  is  reduced  by  the  latter  to  Fe,  the 
other  product  of  the  reaction,  water  (H2O),  being 
absorbed  in  the  calcium  chloride  tube  c.  Care  must  be 
taken  to  volatilize  and  drive  into  c  any  water  which 
condenses  in  the  end  of  the  hard  glass  tube  b. 

After  heating  about  15  minutes,  take  away  the  flame 
and  allow  the  hard  glass  tube  and  contents  to  cool,  the 
current  of  hydrogen  being  continued,  however.  When 
the  tube  is  thoroughly  cooled,  disconnect  the  apparatus 
and  quickly  cap  the  ends  of  c.  Then  carefully  weigh 
b  and  also  c. 

What  has  been  the  decrease  in  the  weight  of  the  Fe2O3? 


10 


EXPERIMENTS  IN  GENERAL  CHEMISTRY 


What  does  this  loss  represent?  What  is  the  increase 
in  the  weight  of  c?  What  does  this  increase  represent? 
From  the  loss  in  weight  of  the  Fe203,  calculate  what 
should  have  been  the  increase  in  the  weight  of  c,  and 
compare  this  figure  with  that  found  experimentally. 
What  is  the  percentage  error?  Was  the  Fe2O3  in  b 
completely  reduced  ?  If  not,  what  per 
cent  of  the  Fe2O3  was  not  reduced  ? 
Arrange  all  data  in  tabular  form 

DETERMINATION  or  THE  EQUIVALENT 
CpO  WEIGHT  OF  MAGNESIUM. 

(Quantitative.) 

15.  First  Method.  By  Means  of 
the  Eudiometer.  Accurately  weigh 
out  0.03  gm.  of  magnesium  (Mg)  rib- 
bon and  introduce  it  into  a  porcelain 
crucible.  Fi^l  the  crucible  with  water 
and  lower  it  into  a  beaker  likewise 
filled  with  water.  Now  fill  the  long 
arm  of  the  eudiometer  with  water  and 
by  covering  the  end  with  the  finger, 
invert  the  tube  in  the  beaker  of  water 
and  clamp  in  position  as  shown  in 
Fig.  5.  The  tube  should  entirely 
cover  the  metal  in  the  crucible  and 
should  be  completely  filled  with  water 
^  below  the  stopcock. 
FIG.  5.  Partially  fill  the  upper  end  of  the 

eudiometer  with  concentrated  HC1. 
Now  carefully  open  the  stopcock  to  allow  a  few  cubic 
centimeters  of  the  acid  to  run  down  into  the  water.  On 


HYDROGEN  II 

account  of  its  greater  specific  gravity,  the  acid  gradu- 
ally settles  to  the  bottom  of  the  tube  and  comes  into 
contact  with  the  magnesium.  The  action  is  slow  at  first 
but  gradually  increases.  Care  must  be  exercised  to  pre- 
vent the  introduction  of  too  much  acid. 

When  the  metal  is  entirely  dissolved,  close  the  end  of 
the  tube  under  the  water  by  means  of  the  finger  and 
transfer  the  tube  to  a  tall  cylinder  filled  with  water. 
Raise  the  tube  so  that  the  surface  of  the  water  within 
is  at  the  same  level  as  the  water  in  the  cylinder,  and 
then  carefully  read  the  volume  of  gas  in  the  eudiometer. 
Also  take  the  temperature  of  the  water  and  note  the 
barometric  pressure  in  the  room.  The  observed  pres- 
sure is  not  exactly  the  same  as  the  pressure  within  the 
tube,  on  account  of  the  water  vapor  in  the  latter.  The 
difference  in  pressure  is  equal  to  the  aqueous  tension  at 
the  observed  temperature. 

The  volume  of  gas  must  now  be  corrected  to  standard 
conditions,  that  is,  to  o°  C.  and  760  mm.  pressure.  Let: 

T  =  observed  temperature, 

r  =  o°  c.  (273°  Ab.), 

P  =  observed  pressure, 
a  =  aqueous  tension  at  T, 
P  —  a  =  actual  pressure  in  the  tube, 
P'  =  760  mm., 
V  =  observed  volume, 
Fo  =  volume  corrected  to  o°  and  760  mm. 

According  to  the  laws  of  Charles  and  Boyle,* 

VPT* 

*-'W 

*  See  Appendix,  page  181. 


12  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

but,  inasmuch  as  the  actual  pressure  within  the  tube  is 
not  P  but  is  P  —  a,  the  expression  becomes: 

V(P-d)Tr 


The  weight  of  i  liter  (1000  cc.)  of  hydrogen  at  stand- 
ard conditions  is  0.0899  gm-  J  hence  the  weight  of  i  cc. 
is  0.0000899  gm->  and  the  weight  of  FO  cc.  may  be 
expressed  by  the  formula: 

Fo  X  0.0000899  gm. 

The  equivalent  weight  of  magnesium  is  the  weight  of 
that  metal  which  is  equivalent  to  i  gm.  of  hydrogen. 
Inasmuch  as  0.03  gm.  of  magnesium  is  equivalent  to 
Fo  X  0.0000899  Sm-  °f  hydrogen,  the  amount  equivalent 
to  i  gm.  of  the  latter  is  found  by  the  proportion: 

(Fo  X  0.0000899)  :  °-°3  ::  *  :  x> 


in  which  x  equals  the  equivalent  weight  of  magnesium. 

1  6.  Repeat  the  above  experiment,  using  aluminum 
(Al)  or  zinc  (Zn)  instead  of  magnesium.  When  per- 
forming the  experiment  with  zinc,  employ  o.i  gm.  of  the 
metal  instead  of  0.03  gm.  as  directed  above.  Why? 

Compare  the  equivalent  weights  of  the  metals  deter- 
mined with  their  respective  atomic  weights.  Does  this 
comparison  bring  out  any  noteworthy  facts?  What 
can  you  judge  as  to  the  valence  of  the  metals  ? 

17.  Second  Method.  Without  Special  Apparatus. 
This  method  is  much  more  crude  than  that  described  in 
Exp.  15,  but  inasmuch  as  the  volume  of  gas  liberated  is 
large,  the  relative  error  is  small,  hence  the  method  yields 
fairly  accurate  results. 

Prepare  an  apparatus  as  shown  in  Fig.  2,  page  5. 


HYDROGEN  13 

Accurately  weigh  out  about  5  gms.  of  pure  zinc  and 
introduce  it  into  the  flask.  Through  the  thistle  tube 
add  sufficient  water  to  cover  the  zinc.  Place  an  inverted 
bottle  filled  with  water  over  the  end  of  the  exit  tube  and 
have  several  other  bottles  filled  with  water  ready  to 
substitute  for  the  first  as  soon  as  it  is  filled  with  gas. 

Now  add  through  the  thistle  tube  a  measured  volume 
of  concentrated  HC1,  10-20  cc.  If  this  does  not  pro- 
duce a  brisk  evolution  of  hydrogen,  add  more  acid,  being 
careful  to  note  the  exact  volume.  Collect  all  the  gas 
coming  from  the  exit  tube  and  use  care  to  prevent  loss 
of  gas  while  changing  the  bottles. 

When  the  zinc  is  entirely  dissolved,  ascertain  the  total 
volume  of  gas  which  has  been  collected.  The  volume 
of  each  bottle  can  readily  be  found  by  filling  the  bottle 
with  water  and  then  measuring  the  latter  by  means  of  a 
graduated  cylinder.  The  total  volume  of  gas  collected 
differs  from  the  exact  volume  of  hydrogen  liberated  only 
by  the  volume  of  the  concentrated  HC1  employed.  The 
volume  of  hydrogen  generated,  F,  is  found,  therefore, 
by  subtracting  the  volume  of  concentrated  HC1  used 
from  the  total  volume  of  gas  collected. 

Read  the  barometer  and  take  the  temperature  of  the 
water  over  which  the  gas  was  collected.  The  volume 
of  hydrogen  must  now  be  reduced  to  standard  condi- 
tions. Ignoring  the  aqueous  tension,  and  making  use 
of  the  formula  given  in  Exp.  15,  we  have 

VPT' 

Vo  =-  W> 

and  the  weight  of  the  hydrogen  equals 
F0  X  0.0000899  gm. 


14  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

The  equivalent  weight  of  zinc  is  then  obtained  by  the 
following  proportion: 

(Fo  X  0.0000899)  \Wt.  ::  i  :  x, 

in  which  Wt.  is  the  weight  of  zinc  used,  and  x  is  the 
equivalent  weight  of  the  metal. 


CHAPTER  III. 

OXYGEN  AND   OZONE. 

OXYGEN  (O;  16). 

1 8.  Preparation.  In  hard  glass  test  tubes  heat  separ- 
ately small  amounts  of  each  of  the  following  substances: 
mercuric  oxide  (HgO),  barium  peroxide  (Ba02),  potas- 
sium chlorate  (KC103),  potassium  dichromate  (K2Cr207), 


FIG.  6. 

potassium  nitrate  (KN03).  In  each  case  test  the  gas 
evolved  by  introducing  a  glowing  splinter  into  the  mouth 
of  the  tube. 

19.   Mix  about  10  gms.  of  manganese  dioxide  (Mn02) 
and  10  gms.  of  potassium  chlorate  (KC103).    Arrange 

is 


l6  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

an  apparatus  as  shown  in  Fig.  6,  and  introduce  the 
mixture  of  MnO2  and  KC103  into  the  test  tube.  Apply 
heat  very  slowly  and  gradually.  After  a  regular  flow 
of  oxygen  comes  from  the  apparatus,  collect  several 
bottles  of  the  gas  by  displacement  of  water.  (These 
bottles  of  oxygen  are  for  use  in  the  following  experiments 
on  the  properties  of  the  element.) 

What  reaction  has  taken  place  in  the  above  prepara- 
tion of  oxygen  ?  Why  is  the  Mn02  mixed  with  the  KClOa  ? 

20.  Introduce    about    2    gms.    of    sodium    peroxide 
(Na2O2)  into  a  test  tube  and  add  about  5  cc.  of  water. 
What   happens?     Test   the   gas   evolved.     Is   the   gas 
oxygen  ? 

21.  Properties.    Dry  a  small  piece  of  phosphorus  (P) 
by  means  of  filter  paper  and  place  it  in  a  deflagrating 
spoon.     Ignite  the  phosphorus  by  touching  it  with  a 
hot  file  and  then  thrust  into  a  bottle  of  oxygen.     What 
causes  the  white  fumes  ? 

When  the  phosphorus  has  stopped  burning,  withdraw 
the  deflagrating  spoon  from  the  bottle  and  introduce 
a  piece  of  moistened  blue  litmus  paper.  Does  the  paper 
change?  If  so,  what  causes  the  change?  (CAUTION! 
The  greatest  care  must  be  exercised  in  experiments  in- 
volving the  use  of  phosphorus.  Phosphorus  is  spontane- 
ously inflammable  in  the  air.  After  finishing  the  above 
experiment  the  iron  deflagrating  spoon  should  be  strongly 
ignited  in  the  Bunsen  burner  flame  to  burn  any  remaining 
traces  of  phosphorus.  Never  handle  phosphorus  with  the 
fingers;  the  heat  of  the  hand  is  sufficient  to  cause  it  to 
ignite.  Phosphorus  burns  are  very  painful.) 

22.  Fill  a  deflagrating  spoon  with  flowers  of  sulphur 
(S).    Ignite  in  the  Bunsen  flame  and  quickly  introduce 


OXYGEN  17 

into  a  bottle  of  oxygen.  When  combustion  is  complete, 
withdraw  the  deflagrating  spoon  and  hold  a  strip  of  wet 
blue  litmus  paper  in  the  mouth  of  the  bottle.  Explain 
fully  what  produces  the  change  in  the  litmus  paper. 

23.  Unravel  one  end  of  a  piece  of  picture  wire  for 
about  i  cm.  Heat  the  unravelled  end  to  redness  and 
quickly  plunge  into  a  bottle  of  oxygen.  Notice  the 
brilliancy  of  the  combustion.  What  are  the  metallic 
looking  globules  formed  ?  Introduce  a  piece  of  moistened 
litmus  paper  into  the  mouth  of  the  bottle.  Is  the  paper 
changed  ? 

Summary.  In  what  respects  does  oxygen  differ  from 
hydrogen  ?  What  test  would  you  use  to  identify  oxygen  ? 
Why  does  a  splinter  burn  more  readily  in  oxygen  than 
in  air?  Should  oxygen,  like  hydrogen,  be  kept  in  in- 
verted bottles?  Why? 

Problems,  (a)  What  volume  of  oxygen  can  be  prepared  by 
igniting  200  gms.  of  mercuric  oxide? 

(b)  The  gas  chamber  in  a  gasometer  is  40  cm.  in  diameter  and 
90  cm.  high.    What  weight  of  a  mixture  of  equal  parts  of  MnO2 
and  KClOa  will  be  necessary  to  generate  enough  oxygen  to  fill  the 
gasometer? 

(c)  What  weight  and  volume  of  oxygen  can  be  obtained  by  the 
complete  electrolysis  of  i   liter  of  water?     What  weight  and 
volume  of  hydrogen  will  be  produced  at  the  same  time? 

DETERMINATION  OF  THE  WEIGHT  OF  A  LITER  OF  OXYGEN. 
(Quantitative.) 

24.  Arrange  the  apparatus  shown  in  Fig.  7.  The 
bottle  should  be  of  about  2  liters  capacity;  all  stoppers 
should  be  of  rubber.  Completely  fill  the  bottle  a  with 
water;  insert  the  stopper  carrying  the  two  glass  tubes 


i8 


EXPERIMENTS   IN   GENERAL  CHEMISTRY 


as  shown.  Water  will  be  forced  up  through  tube  b 
completely  filling  it.  Now  clamp  this  tube  by  means 
of  a  pinchcock  as  shown.  Insert  tube  b  in  beaker  c 
which  should  be  empty. 

Remove  the  test  tube  d  and  fill  it  about  one  third  full 
of  a  mixture  of  equal  parts  of  KC103  and  MnO2,  both 
of  which  have  been  previously  dried  by  gently  warming 


FIG.  7. 

in  a  porcelain  evaporating  dish.  Weigh  the  tube  with 
the  mixture  and  let  w'  represent  this  weight.  Connect 
the  test  tube  with  the  apparatus,  and  then  open  the 
pinchcock  on  b.  If  all  connections  are  tight,  no  water 
will  run  into  c. 

Apply  heat  slowly  and  gently  to  the  end  of  the  test 
tube  containing  the  mixture  of  KC1O3  and  Mn02.  If 
too  strong  a  heat  is  applied,  oxygen  will  be  evolved  too 
rapidly  and  may  result  in  loss.  The  gas  should  be 
evolved  slowly  and  steadily,  The  oxygen  which  col- 


OXYGEN  19 

lects  in  bottle  a  displaces  its  own  volume  of  water, 
forcing  it  over  into  beaker  c. 

The  heating  is  continued  until  bottle  a  is  about  half 
full  of  oxygen.  Then  remove  the  source  of  heat  and 
allow  the  apparatus  to  cool.  When  cool,  raise  beaker  c 
until  the  surface  of  the  water  contained  therein  and 
the  surface  of  the  water  in  a  are  at  the  same  level. 
While  holding  the  beaker  in  this  position,  tighten  the 
pinchcock  on  tube  b  and  then  remove  the  latter  from  c. 

Measure  the  volume  of  the  water  in  c.  This  is  the 
same  as  the  volume  of  oxygen  in  a.  Take  the  temperature 
of  the  water  in  c\  the  oxygen  in  a  is  at  the  same  temper- 
ature. Note  barometric  pressure  in  the  room. 

Carefully  remove  the  test  tube  d  from  its  stopper  and 
weigh  accurately.  Let  this  weight  be  represented 
by  w"  '.  Then  w'  —  w"  equals  the  weight  of  the  oxygen 
evolved. 

The  volume,  V,  must  now  be  corrected  for  temperature 
and  pressure.  This  is  done  in  exactly  the  same  manner 
as  described  in  Exp.  15,  by  the  formula: 

V(P  -  a)T' 

Vo=    ~- 


If  FO  cc.  of  oxygen  under  standard  conditions  of 
temperature  and  pressure  weigh  wr  —  w"  gms.,  the 
weight  of  i  liter  (1000  cc.)  of  oxygen  can  be  obtained 
from  the  proportion: 

FO  :  (wr  —  w")  :  :  1000  :  x, 

in  which  x  equals  the  weight  of  a  liter  of  oxygen  at  the 
standard  conditions  of  temperature  and  pressure. 

Calculate  the  theoretical  weight  of  a  liter  of  oxygen 


2O  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

and  compare  with  the  result  obtained  experimentally. 
What  is  the  percentage  error  ?  To  what  do  you  attribute 
the  error  ? 

In  this  experiment  is  it  necessary  to  drive  all  the 
available  oxygen  from  the  mixture  in  the  tube  ?  Why  ? 

VERIFICATION  OF  THE  LAW  OF  DEFINITE 
PROPORTIONS. 

25.  Part  I.  Carefully  weigh  a  clean,  dry  porcelain 
crucible  and  cover.  By  means  of  fine  sand-paper 
thoroughly  clean  a  piece  of  magnesium  ribbon  about 
3  feet  long.  Wipe  the  ribbon  with  a  towel  to  remove 
particles  of  dust.  Twist  the  ribbon  into  a  coil  and 
press  down  firmly  into  the  crucible.  Replace  the  cover 
and  again  weigh.  The  difference  in  the  two  weighings 
is  the  weight  of  magnesium  taken. 

Place  the  crucible  on  a  pipestem  triangle  on  a  ring 
stand  and  heat  gently  with  a  small  flame.  The  burner 
should  be  held  in  the  hand.  With  the  other  hand,  by 
means  of  a  pair  of  clean  crucible  tongs  or  steel  pincers, 
hold  the  cover  of  the  crucible  a  little  above  the  crucible 
in  order  that  air  may  enter  the  latter.  If  the  magnesium 
ribbon  takes  fire,  instantly  cover  the  crucible  and  with- 
draw the  flame  until  the  burning  within  the  crucible 
ceases.  Then  again  apply  heat,  holding  the  crucible 
cover  as  before.  Great  care  must  be  taken  to  prevent 
loss  of  any  of  the  white  powder  which  clings  to  the  under 
side  of  the  cover. 

When  oxidation  seems  to  be  complete,  that  is,  when 
the  magnesium  ribbon  no  longer  takes  fire,  entirely 
remove  the  cover,  being  careful  to  lose  none  of  the  white 


OXYGEN  21 

oxide  which  clings  thereto.  Place  the  crucible  in  an 
inclined  position  on  the  triangle  and  heat  strongly  for  a 
few  minutes.  Then  allow  the  crucible  to  cool,  replace 
the  cover  and  again  weigh. 

In  order  to  make  sure  that  oxidation  is  complete,  the 
crucible  without  the  cover  should  again  be  placed  in  an 
inclined  position  on  the  triangle  and  strongly  heated  for 
a  few  minutes.  Allow  to  cool;  cover  and  weigh.  The 
two  successive  weighings  should  be  identical. 

What  is  the  weight  of  the  contents  of  the  crucible? 
How  much  has  the  magnesium  gained  in  weight  ?  Cal- 
culate the  percentage  gain  in  weight. 

Part  II.  Carefully  weigh  a  small  clean  porcelain 
evaporating  dish.  Clean  about  3  feet  of  magnesium 
ribbon  as  described  in  Part  I,  twist  into  a  coil,  introduce 
into  the  dish  and  carefully  weigh.  The  difference  in  the 
two  weighings  equals  the  amount  of  magnesium  taken. 

To  the  contents  of  the  dish  add  enough  distilled  water 
to  cover  the  magnesium.  Then  add  pure  concentrated 
HNOs,  a  few  drops  at  a  time,  quickly  covering  the  dish 
with  a  watch  glass  after  each  addition.  (Why  ?)  When 
the  metal  is  entirely  dissolved,  place  the  dish  on  a  water 
bath  and  evaporate  to  dryness.  Transfer  the  dish  to  a 
triangle  on  a  ring  stand  and  apply  heat,  very  carefully  at 
first,  to  prevent  loss,  and  then  more  strongly  until  brown 
fumes  are  no  longer  evolved.  Allow  the  dish  to  cool 
and  then  weigh. 

How  much  has  the  magnesium  gained  in  weight? 
What  is  its  percentage  gain  in  weight  ? 

Compare  the  percentage  gain  in  weight  found  in  Part 
II  with  that  found  in  Part  I.  What  conclusion  can  you 
draw  from  these  two  experiments  ? 


22  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

OZONE. 

26.  In   a  good-sized  beaker  place  a   few  pieces  of 
yellow  phosphorus.      (CAUTION!    See  page  16.)     Add 
enough  water  to  partially  cover  the  phosphorus.     Cover 
the  beaker  with  a  watch  glass  and  allow  to  stand  for 
about  10  minutes. 

Remove  the  watch  glass  and  notice  the  odor  of  the  air 
in  the  beaker.  Test  the  air  for  ozone  by  introducing  a 
piece  of  filter  paper  moistened  with  starch  paste  con- 
taining a  little  potassium  iodide  (KI).  What  do  you 
notice?  Explain  all  the  chemical  changes  which  have 
taken  place  in  the  test  solution. 

27.  Pour  a  few  drops  of  ether  into  a  beaker  or  a  wide 
mouth  bottle  and  quickly  cover.     (CAUTION!    Ether  is 
very  inflammable  and  should  never  be  poured  from  a  bottle 
when   a  flame   is   near.)     Keeping    the   bottle    tightly 
covered,  shake  it  to  completely  cover  the  sides  with  a  film 
of  ether. 

Heat  the  end  of  a  glass  rod  almost  to  redness  and 
plunge  it  into  the  bottle.  If  the  rod  is  too  hot,  an 
explosion  will  occur.  Test  the  atmosphere  in  the 
beaker  with  filter  paper  saturated  with  the  Kl-starch 
paste  test  solution.  Also  note  the  odor  coming  from 
the  bottle. 

Summary.  How  can  ozone  be  prepared  in  larger 
quantities?  Does  ozone  ever  occur  in  the  air?  To 
what  cause  may  this  be  due?  Does  oxygen  affect 
Kl-starch  paste?  Why  should  ozone  be  more  active 
than  oxygen?  Mention  one  other  good  test  for  ozone. 


CHAPTER  IV. 

WATER  AND  HYDROGEN  PEROXIDE. 
WATER. 

28.  Distilled  Water.    Arrange  an  apparatus  consisting 
of  a  condenser  and  a  distilling  flask  fitted  with  a  ther- 


mometer as  shown  in  Fig.  8.  Partially  fill  the  flask 
with  water  and  heat  to  boiling.  Collect  the  distilled 
water  which  runs  from  the  condenser.  Continue  the 
distillation  until  about  150  cc.  are  obtained.  Notice 
the  temperature  during  distillation. 

What  is  meant  by  distillation  ?  Why  is  distilled  water 
purer  than  the  ordinary  city  service  water? 

29.  On  a  clean  watch  glass  evaporate  to  dryness  a 
few  cubic  centimeters  of  the  ordinary  city  service  water, 
or,  better  still,  a  few  cubic  centimeters  of  river  water. 

23 


24  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

Notice  the  amount  of  residue.  (In  making  this  evapo- 
ration, place  the  watch  glass  on  a  wire  gauze  on  a  ring 
stand  and  heat  gently  by  means  of  a  small  flame.)  On 
a  second  watch  glass  evaporate  a  similar  volume  of  dis- 
tilled water  and  compare  the  amount  of  residue  with 
that  previously  obtained.  What  is  your  deduction  as 
to  the  relative  purity  of  the  two  samples  of  water  ? 

Solubility  in  Water. 

30.  Solids.    Test  the  solubility  of  a  number  of  salts 
and  arrange  the  results  in  two  columns  headed  respec- 
tively " Soluble"  and  "Insoluble." 

To  make  a  solubility  test,  introduce  a  small  piece  of 
the  material  to  be  tested,  not  larger  than  a  grain  of 
wheat,  into  a  test  tube  half  full  of  distilled  water.  Close 
the  tube  with  the  thumb  and  shake  vigorously.  If  the 
substance  does  not  seem  to  dissolve,  heat  to  boiling  for 
a  few  moments  and  then  allow  the  tube  to  stand  several 
minutes. 

Any  of  the  following  substances  may  be  used  in  this 
experiment:  copper  sulphate  (CuSO^,  ferrous  sulphate 
(FeSO4),  mercurous  chloride  (HgCl),  potassium  dichro- 
mate  (K2Cr2O7),  potassium  chloride  (KC1),  calcium  car- 
bonate (CaCO3),  magnesium  sulphate  (MgSOJ,  borax 
.(Na2B4O7)  and  ammonium  chloride  (NHaCl). 

Compare  the  results  obtained  with  the  table  of  solu- 
bilities on  page  200.  Are  there  any  discrepancies?  If 
so,  to  what  are  the  discrepancies  due  ? 

31.  Liquids.     Test  the  solubility  of  several  liquids  in 
water  by  adding  about  3  cc.  of  each  to  a  test  tube 
half  full  of  distilled  water.     In  each  case  if  the  liquid 
does  not  dissolve  immediately,  shake  the  tube  gently 


WATER  25 

and  then  allow  to  stand  for  a  moment.    Arrange  the 
data  obtained  as  in  the  previous  experiment. 

The  following  compounds,  which  are  liquids,  may  be 
used  in  this  experiment:  ether,  alcohol,  kerosene, 
chloroform,  glycerine  and  carbon  disulphide. 

DETERMINATION  OF  THE   SOLUBILITY  OF   SODIUM 

CHLORIDE. 

(Quantitative.) 

32.  Make  a  saturated  solution  of  sodium  chloride 
(NaCl)  by  treating  about  100  gms.  of  the  salt  with 
200  cc.  of  water  in  a  small  flask.  Allow  to  stand,  with 
occasional  shaking,  for  about  half  an  hour.  While  the 
mixture  is  standing,  accurately  weigh  a  small  porcelain 
evaporating  dish. 

When  the  solution  is  saturated,  i.e.,  when  no  more 
salt  will  dissolve,  filter  a  portion  of  the  solution  through 
a  plaited  filter,  catching  the  filtrate  in  the  evaporat- 
ing dish.  The  dish  should  be  about  half  full  of  the 
solution.  Take  the  temperature  of  the  solution  in  the 
dish.  Carefully  weigh  dish  and  solution  and  deduct 
the  weight  of  the  empty  dish,  thus  arriving  at  the  exact 
weight  of  the  salt  solution.  Let  this  weight  be  repre- 
sented by  wr. 

Evaporate  the  solution  on  a  water  bath,  allowing  the 
dish  to  remain  on  the  bath  until  the  salt  is  perfectly 
dry.  When  cool,  again  weigh  dish  and  contents.  Sub- 
tract the  weight  of  the  empty  dish  in  order  to  find  the 
exact  weight  of  the  residue  —  sodium  chloride.  Repre- 
sent the  weight  of  sodium  chloride  by  w" . 

The  actual  weight  of  the  water  in  which  the  w" 
grams  of  salt  were  dissolved  is  equal  to  w1  —  w" \  The 


26  EXPERIMENTS   IN  GENERAL  CHEMISTRY 

solubility  of  a  substance  at  any  given  temperature  is  the 
number  of  grams  of  the  substance  that  will  dissolve  in 
100  gms.  of  water  at  that  temperature.  If  w"  grams  of 
salt  will  dissolve  in  wr  —  w"  grams  of  water  at  temper- 
ature /,  the  number  of  grams  that  will  dissolve  in  100  gms. 
of  water  at  that  temperature  can  be  calculated  by  the 
proportion : 

(wf  —  w'f)  :  w"  ::  100  :  x, 

in  which  x  equals  the  solubility  of  sodium  chloride  at 
temperature  t. 

Compare  the  result  obtained  with  a  table  of  the  solu- 
bility of  sodium  chloride  and  report  the  percentage  error. 
Draw  a  curve  representing  the  solubility  of  sodium  chlo- 
ride from  o°  to  1 00°  C. 

EXAMINATION  OF  WATER  FOR  IMPURITIES. 

(Qualitative.) 

33.  Each  of  the  following  tests  should  be  made  with 
distilled  water  and  with  samples  of  several  other  waters 
from  as  widely  differing  sources  as  possible.  City  ser- 
vice, river,  canal,  ocean,  lake,  well  or  rain  water  may  be 
used.  Chemically  pure  reagents  must  be  used  in  making 
these  tests.  Compare  the  results  obtained  in  each  test. 

Lime.  To  each  of  the  samples  of  water  in  test  tubes, 
add  a  few  drops-  of  ammonium  hydroxide  and  a  few 
drops  of  ammonium  oxalate  ((NH4)2C204).  Heat  each 
tube  to  boiling  and  then  allow  to  stand  and  settle. 
The  precipitate  of  calcium  oxalate  (CaQjOj  shows  the 
presence  of  calcium  salts  in  the  water.  Calcium  salts 
are  always  reported  as  lime. 

Sulphates.  Test  the  several  samples  of  water  for  sul- 
phates by  adding  to  each  a  drop  of  HC1  and  a  few  drops 


WATER  27 

of  barium  chloride  (BaCl2)  solution.  Heat  to  boiling 
and  then  allow  to  stand  and  settle.  The  white  precipi- 
tate is  barium  sulphate,  BaSO4. 

Chlorides.  Make  the  test  for  chlorides  by  adding  to 
each  sample  a  drop  of  HN03  and  a  few  drops  of  silver 
nitrate  (AgN03)  solution. 

Ammonia.  The  test  solution  for  ammonia  is  Nessler's 
reagent,  which  is  a  mixture  of  potassium  mercuric  iodide 
(HgI2.2  KI)  and  potassium  hydroxide  (KOH)  solutions. 
With  waters  containing  ammonia  or  ammonium  salts, 
the  reagent  produces  a  yellowish  brown  color,  the  depth 
of  color  being  indicative  of  the  amount  of  ammonia 
present.  The  color  can  be  most  easily  judged  by  look- 
ing down  through  the  tube  at  a  piece  of  white  paper. 

Test  the  several  samples  of  water  for  ammonia  by 
adding  to  equal  volumes  of  each  i  cc.  of  Nessler's  re- 
agent. Allow  the  tubes  to  stand  for  two  or  three  min- 
utes before  making  the  comparison. 

34.  Hardness  of  Water.    Add  about  5  cc.  of  soap 
solution  to  100  cc.  of  distilled  water  in  a  clean  500  cc. 
flask.     Shake  the  flask  for  several  minutes,  and  notice 
the    sound    which    is    produced.     Does    a    permanent 
lather  form?     The  formation  of  a  lather  and  the  pro- 
duction of  very  little  noise  when  the  flask  is  shaken  indi- 
cate that  the  water  is  soft. 

« 

Test  the  hardness  of  several  other  samples  of  water 
by  adding  a  little  soap  solution  to  100  cc.  of  each  as 
described  above.  Record  all  results. 

35.  Repeat  the  test  with  one  of  the  samples  of  hard 
water,  using,  instead  of  5  cc.  of  the  soap  solution,  at 
least  10  or  15  cc.     Shake  as  before.     Do  you  get  the 
same  results  as  with  this  water  in  Exp.  34?    What  is 


28 


EXPERIMENTS   IN   GENERAL  CHEMISTRY 


the  reason?    What  can  be  added  to  a  hard  'water  to 

make  it  soft  ? 
36.   Repeat  the  experiment  with  distilled  water  and 

5  cc.  of  soap  solution,  first,  however,  adding  to  the  dis- 
tilled water  10  or  15  cc.  of  a  solution 
of  any  calcium  or  magnesium  salt. 
Compare  with  the  results  obtained 
from  Exp.  34. 

What  has  caused  the  change  ?  Is 
the  water  now  hard  or  soft  ?  If  it  is 
hard,  how  can  it  be  softened  ? 

DETERMINATION  or  THE  HARD- 
NESS OF  WATER. 
(Quantitative.) 

37.  Obtain  a  supply  (about  50  cc.) 
of  "standard  soap  solution"  from  the 
stock  bottle  on  the  reagent  shelves. 
Rinse  out  a  burette  three  times  with 
small  amounts  of  the  soap  solution 
and  then  fill  the  burette  with  it 
(Fig.  9). 

Into  a  clean  glass-stoppered  bottle 
introduce  exactly  100  cc.  of  the  ordi- 
nary city  service  water.  Note  the 
level  of  the  solution  in  the  burette, 
then  allow  the  solution  to  run,  a  drop 
at  a  time,  into  the  bottle  of  water, 
shaking  the  bottle  vigorously  after 


FIG.  9. 


each  addition.  Continue  to  add  soap  solution  in  this 
way  until  a  drop  finally  produces  a  permanent  lather. 
Read  the  burette  again  and  note  the  number  of  cubic 


WATER  29 

centimeters  of  soap  solution  required  to  produce  the 
lather,  i.e.,  to  discharge  the  hardness  of  the  water. 
The  standard  soap  solution  is  so  made  up  that  i  cc.  is 
equivalent  to  i  part  of  calcium  carbonate,  i.e.,  hardness, 
in  100,000  parts  of  water.  Express  your  results  in  parts 
of  hardness  per  100,000. 

This  experiment  should  be  repeated  two  or  three 
times  and  the  results  averaged.  All  results,  however, 
should  appear  in  the  report. 

In  a  similar  manner  determine  the  hardness  of  one  or 
two  other  samples  of  water,  as  river  water,  or  well  water. 

38.  Water  of   Crystallization.    Notice   the   color   of 
crystals  of  copper  sulphate  (CuSO4).     Powder  several 
crystals  and  notice  the  color  of  the  powder.     Introduce 
about  10  gms.  of  the  powder  into  a  porcelain  evaporat- 
ing dish  and  heat  gently  until  the  powder  is  white. 
What  kind  of  a  change  has  taken  place?     Allow  the 
dish  and  powder  to  cool;    then  add  a  few  cubic  centi- 
meters   of   water.     Notice    the    immediate    change    in 
appearance.     Explain  all  changes. 

39.  Place  a  crystal  of  copper  sulphate  in  a  test  tube 
and  apply  heat.     Observe  that  the  water  driven  out  of 
the  crystal  condenses  at  the  outer  end  of  the  tube.     In 
like  manner,  in  separate  test  tubes,  heat  crystals  of  the 
following   substances:     magnesium   sulphate    (MgSGU), 
potassium    dichromate    (K2Cr207),   potassium    chloride 
(KC1),  ferrous  sulphate  (FeSO4)  and  borax  (Na2B4O7). 
Judging  by  the  amount  of  water  in  each  case,  which  of 
these  salts  do  you  conclude  contain  water  of  crystalliza- 
tion? 

Compare  the  results  of  this  experiment  with  those 
obtained  in  Exp.  30.  Do  you  conclude  that  all  salts 


30  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

having  water  of  crystallization  are  soluble  in  water? 
Do  you  conclude  that  all  salts  having  no  water  of  crystal- 
lization are  insoluble  ? 

40.  Obtain  about  one  third  of  a  test  tube  full  of 
sodium  acetate   (NaC2H302)   from  the  reagent  bottle. 
Add  about  i  cc.  of  water  and  heat  gently  until  all  of  the 
crystals  have  dissolved  and  the  test  tube  contains  only 
a  perfectly  clear  solution.    Place  a  plug  of  cotton  in  the 
mouth  of  the  tube  and  allow  the  latter  to  stand  in  the 
test  tube  rack  undisturbed  until  perfectly  cool.     Notice 
that  the  tube  still  contains  a  perfectly  clear  solution. 

Remove  the  plug  of  cotton  and  drop  into  the  tube  a 
small  crystal  of  sodium  acetate.  What  phenomenon 
do  you  observe?  Does  the  temperature  of  the  tube 
change?  Can  you  offer  an  explanation  of  the  phenom- 
enon? 

41.  Efflorescence.     On  separate  watch  glasses,  expose 
to  the  air  for  several  days  a  few  large  clear  crystals 
of  each  of  the  following  substances:   ferrous  sulphate 
(FeSO4),  sodium  sulphate  (Na2SO4)  and  zinc  sulphate 
(ZnSO4).     What  change  takes  place  during  the  time 
that  they  are  exposed?    Why  does  this  change  occur? 
Mention  one  or  two  salts  which  are  not  efflorescent. 

42.  Deliquescence  and  Hygroscopicity.     On  separate 
watch  glasses  expose  to  the  air,  for  a  day  or  two,  small 
pieces  of  each  of  the  following  substances:  zinc  chloride 
(ZnCl2),  calcium    chloride    (CaCl2),  sodium    hydroxide 
(NaOH)    and    phosphorus    pentoxide    (P2O5).       What 
change  takes  place  in  the  appearance  of  these  substances  ? 
What  is  the  difference  between  the  terms  "deliquescent" 
and  " hygroscopic"?     Which  of  the  above-mentioned 
substances  come  in  each  class  ? 


WATER  31 

43.  Water  of  Decrepitation.    Heat  two  or  three  large 
crystals  of  sodium  chloride  (NaCl)  in  a  test  tube.     Ex- 
plain the  phenomenon  observed. 

In  a  porcelain  mortar  grind  a  few  crystals  of  NaCl 
to  a  fine  powder.  Introduce  a  portion  of  this  powder 
into  a  test  tube  and  heat.  Why  does  the  result  differ 
from  that  obtained  when  large  crystals  were  employed? 

DETERMINATION  OF  THE  NUMBER  or  MOLECULES  or 
WATER  OF  CRYSTALLIZATION  IN  GYPSUM. 

(Quantitative.) 

44.  Carefully  weigh  a  clean  dry  porcelain  crucible 
without  the  cover.     About  half  fill  the  crucible  with 
powdered  gypsum   and   again  weigh  accurately.     The 
difference  in   the   two   weighings  gives   the   weight   of 
gypsum  taken. 

Place  the  crucible  on  a  pipestem  triangle  on  a  ring 
stand  and  apply  heat,  very  gently  for  a  time,  and  then 
gradually  stronger  until  the  full  force  of  the  Bunsen 
burner  is  employed.  Continue  the  strong  heat  for 
about  5  minutes. 

Allow  the  crucible  to  cool  to  the  room  temperature; 
then  weigh.  Again  place  the  crucible  on  the  triangle 
and  heat  with  the  full  force  of  the  burner  for  5  minutes; 
allow  to  cool  and  weigh.  This  process  of  heating  should 
be  continued  until  two  successive  weighings  are  identical. 

The  loss  in  weight  is  due  to  the  water  of  crystalliza- 
tion which  has  been  expelled.  Calculate  the  percentage 
of  water  of  crystallization  which  the  gypsum  contained. 
What  is  the  percentage  of  residue,  i.e.,  anhydrous  cal- 
cium sulphate? 


32      EXPERIMENTS  IN  GENERAL  CHEMISTRY 

In  order  to  calculate  the  number  of  molecules  of  water 
of  crystallization  which  the  gypsum  originally  contained, 
divide  the  percentage  of  water  of  crystallization  found 
by  the  molecular  weight  of  water.  Let  this  result  be 
represented  by  a.  Likewise  divide  the  percentage  of 
residue  (anhydrous  calcium  sulphate)  by  the  molecular 
weight  of  calcium  sulphate.  Call  this  b.  The  number 
of  times  which  a  is  greater  than  b  equals  the  number  of 
molecules  of  water  of  crystallization  which  the  gypsum 
originally  contained. 


DETERMINATION  OF  THE  SPECIFIC  GRAVITY  OF 
SOLIDS  HEAVIER  THAN  WATER. 

45.    Select  a  suitable  piece  of  the  solid  to  be  deter- 
mined and  after  thoroughly  cleaning,  ascertain  its  exact 


FIG.  10. 


weight  (W).  Then  place  a  small  bench  over  the  pan  of 
the  balance  in  such  a  way  that  it  does  not  touch  the 
latter.  On  the  bench  place  a  beaker  of  distilled  water 
large  enough  to  accommodate  the  sample.  By  means 


WATER  33 

of  a  fine  silk  thread  suspend  the  sample  from  the  arm 
of  the  balance  so  that  the  sample  hangs  in  the  water  and 
is  completely  submerged.  Ascertain  the  weight  of  the 
sample  in  water  by  carefully  adding  weights  to  the  other 
pan  of  the  balance.  Let  the  weight  in  water  be  repre- 
sented by  W.  (Fig.  lo.) 

W  is  less  than  W  because  a  substance  immersed  in 
water  is  buoyed  up  by  the  water.  The  loss  of  weight  in 
water  is  equal  to  the  weight  of  the  volume  of  water  displaced, 
and  from  this,  the  loss  of  weight  in  water  is  equal  to  the 
weight  of  an  equal  volume  of  water.  But  the  specific  grav- 
ity is  equal  to  the  weight  of  the  substance  divided  by  the 
weight  of  an  equal  volume  of  water;  therefore, 

W 

Specific  Gravity  =  w_w,  • 


HYDROGEN    PEROXIDE  (H2O2). 

46.  To  about  5  gms.  of  barium  peroxide  (BaO2)  con- 
tained in  a  beaker  add  enough  water  to  form  a  paste. 
Cool  the  mixture  by  adding  a  little  snow  or  ice.     Then 
add  about  25  cc.  of  cold  dilute  sulphuric  acid.     Allow 
to  stand  for  a  few  moments  so  that  the  barium  sulphate 
will  settle.     Pour  off  as  much  as  possible  of  the  super- 
natant liquid  and  filter  it  through  a  double  filter. 

This  liquid  is  a  solution  of  hydrogen  peroxide  (H2O2) 
in  water.     It  is  to  be  used  in  the  following  experiments. 

47.  To    a  portion  of  the  solution  add  a  little  KI- 
starch  paste  or  merely  introduce  a  piece  of  filter  paper 
which  has  been  moistened  with  the  test  solution.    What 
reaction  occurs  ?    What  other  substance  have  we  studied 
which  gives  the  same  color  with  this  test  solution  ? 


34  EXPERIMENTS   IN    GENERAL   CHEMISTRY 

48.  To  another  portion  of  the  solution  add  a  solution 
of  potassium  permanganate  (KMnC^),  drop  by  drop. 
Is  there  a  gas  evolved  ?     What  is  the  gas  ? 

49.  Mix  a  few  drops  of  dilute  sulphuric  acid  and  a 
few  drops  of  potassium  dichromate  solution  in  a  test 
tube  and  to  the  mixture  add  about  a  half  inch  layer  of 
ether.     (CAUTION!     See  page  22,  Exp.  27.)     Now  add 
several  cubic  centimeters  of  hydrogen  peroxide  solution. 
This  is  a  good  test  for  hydrogen  peroxide  and  for  chro- 
mium.    What  color  is  produced  in  the  ether  ? 

50.  To  a  little  powdered  Mn02  in  a  test  tube  add  a 
few  cubic  centimeters  of  hydrogen  peroxide   solution 
taken  from  the  bottle  on  the  side  shelf.     (This  solution 
is  probably  much  stronger  than  the  solution  made  in 
Exp.  46.)     Test  the  gas  evolved.     Explain  the  action. 

51.  Test  the  action  of  H2O2  on  a  solution  of  titanium 
sulphate.     The  color  produced  is  characteristic  and  the 
intensity  depends  upon  the  strength  of  the  titanium 
solution. 

Summary.  State,  in  general  terms,  the  method  for 
the  preparation  of  hydrogen  peroxide.  What  are  two 
important  uses  of  hydrogen  peroxide?  Mention  some 
of  the  common  names  under  which  hydrogen  peroxide  is 
sold  commercially.  'What  is  the  strength  of  commercial 
hydrogen  peroxide  ? 

Problems,  (a)  To  prepare  15  liters  of  2%  H2O2  solution,  what 
weight  of  BaO2  would  be  needed?  What  weight  of  25%  H2SO4 
would  be  required? 

(b)  If  i  liter  of  water  at  4°  C.  is  converted  into  the  form  of 
steam  at  102°,  what  will  be  the  volume  of  the  steam? 

(c)  In  dehydrating  i  ton  of  crystallized  sodium  carbonate  by 
the  aid  of  heat,  what  weight  of  water  would  be  driven  out? 


CHAPTER  V. 

THE  HALOGENS. 

CHLORINE  (Cl;  35). 

52.  Preparation.      In    a    test    tube    treat    about    a 
gram  of  manganese  dioxide  (Mn02)  with  a  few  cubic 
centimeters     of     concentrated     HC1.      Warm    gently. 
Notice  the  yellowish  green  gas  evolved.    Has  it  any 
odor? 

53.  In  separate  test  tubes  try  the  action  of  concen- 
trated HC1  on  small  portions  of  each  of  the  following 
substances:   potassium  chlorate  (KClOs),  barium  perox- 
ide   (Ba02),     potassium    dichromate    (K2Cr2O7),    lead 
dioxide  (Pb02)  and   calcium  hypochlorite    (CaC^O). 

54.  Make  a  mixture  of  about  i  gm.  each  of  Mn02  and 
NaCl.     Introduce  the  mixture  into  a  test  tube  and  treat 
with  H2S04  which  has  previously  been  diluted  with  an 
equal  volume  of  water.     (Pour  the  acid  into  the  water.) 
Warm  the  mixture  and  notice  the  gas  evolved.     Can 
you  explain  the  reaction? 

In  the  above  methods  of  preparation  of  chlorine, 
what  general  principle  is  involved?  Mention  another 
very  important  commercial  method  for  the  preparation 
of  chlorine. 

55.  Arrange  an  apparatus  as  shown  in  Fig.  n.     The 
flask  should  be  of  about  500  cc.  capacity.     The  delivery 
tube  should  extend  to  the  bottom  of  the  bottle.      Place 

35 


30  EXPERIMENTS   IN  GENERAL  CHEMISTRY 

25  or  30  gms.  of  finely  granulated  MnC>2  in  the  flask  and 
through  the  thistle  tube  add  about  100  cc.  of  concen- 
trated HC1.  Agitate  the  flask  to  cause  the  acid  and 
oxide  to  mix  thoroughly. 

Warm  the  flask  gently  and  collect  the  chlorine  evolved 
in  bottles  as  shown  in  the  drawing,  using  a  cardboard 


FIG.  ii. 

or  paper  cover  through  which  the  delivery  tube  passes. 
As  the  bottles  are  filled  they  should  be  covered  with 
glass  plates  on  which  vaseline  has  been  smeared.  Collect 
five  or  six  bottles  of  the  gas.  (They  are  to  be  used 
in  the  following  experiments  on  the  " properties"  of 
chlorine.) 

56.  Chlorine  Water.  Prepare  about  200  cc.  Of  chlor- 
ine water  by  passing  chlorine  gas  through  that  volume 
of  cold  water  contained  in  a  small  flask.  To  hasten 


CHLORINE  37 

the  absorption  of  the  gas,  occasionally  shake  the  flask 
containing  the  water.  Label  this  solution  and  reserve  it 
for  later  experiments. 

57.  Properties  of  Chlorine.    Into  one  of  the  bottles  of 
chlorine  drop  a  little  powdered  antimony  (Sb).    Note 
all   phenomena   observed.    What   compound   or   com- 
pounds are  formed? 

58.  In  the  Bunsen  flame  heat  to  redness  a  thin  strip 
of  copper  (Cu)  foil  and  quickly  plunge  into  a  bottle 
of  chlorine.     Notice  the  products  formed.     Ascertain 
if  they  will  dissolve  in  water  and,  if  so,  what  color  is 
produced. 

59.  Introduce  into   a   third   bottle   of   chlorine,  by 
means  of  a  deflagrating  spoon,  a  small  piece  of  yellow 
phosphorus  which  has  been  dried  by  pressing  gently 
between  pieces  of  filter  paper.     (CAUTION!    See  page  16, 
Exp.  21.) 

60.  Saturate  a  piece  of  filter  paper  with  turpentine 
(Ci0H16),  and  plunge  into  a  bottle  of  chlorine.     Note  the 
flame  and  the  fumes.    Explain  what  has  happened.    Bring 
a  piece  of  wet  blue  litmus  paper  to  the  mouth  of  the 
bottle. 

61.  Moisten  a  piece  of  colored  calico  and  suspend  it  in 
a  bottle  of  chlorine.    Allow  to  stand  for  15  or  20  minutes. 
Is  the  color  of  the  cloth  changed  ? 

Introduce  a  colored  flower  or  a  few  blades  of  grass 
into  a  bottle  of  chlorine  and  observe  the  change  which 
takes  place  after  a  few  moments. 

62.  Add  a  little  chlorine  water  to  a  few  cubic  cen- 
timeters   of   indigo    solution    in    a    test    tube.     What 
change  is  noticed?    What  name  is  applied  to  such  a 
change  ? 


38  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

In  like  manner,  in  separate  test  tubes,  add  chlorine 
water  to  solutions  of:  cochineal,  copper  sulphate 
(CuSO^,  litmus  and  potassium  dichromate  (K2Cr207). 
Why  are  some  of  these  solutions  bleached  and  others 
not? 

In  a  beaker  of  chlorine  water  immerse  a  strip  of 
colored  calico  for  a  few  moments.  Does  the  chlorine 
water  affect  the  calico  in  the  same  manner  as  the  chlor- 
ine gas?  (Compare  with  Exp.  61.) 

Summary.  In  what  respect  is  bleaching  by  chlorine 
similar  to  bleaching  by  hydrogen  peroxide?  If  the  red 
calico  and  the  chlorine  gas  had  both  been  perfectly  dry, 
would  the  cloth  have  been  bleached? 

Explain  in  a  general  way  how  chlorine  bleaches. 

Problems,  (a)  What  volume  of  chlorine  at  20°  and  740  mm. 
pressure  can  be  prepared  by  the  action  of  an  excess  of  HC1  on 
100  gms.  of  K2Cr2O7? 

(b)  In  order  to  prepare  1766  liters  of  chlorine  at  18°  and  755  mm. 
pressure,  what  weight  of  83%  Mn02  will  be  necessary?  What 
weight  of  32%  HC1  will  be  needed?  What  volume  will  this 
amount  of  HC1  have? 


DETERMINATION  OF  THE  WEIGHT  OF  A  LITER 
OF  CHLORINE. 

(Quantitative.) 

63.  Apparatus.  Clean,  dry  flask  of  about  500  cc.  ca- 
pacity, with  tightly  fitting  cork;  thermometer;  balance; 
chlorine  generator  as  shown  in  Fig.  12,  fitted  with  wash 
bottle  a  containing  H2O  and  drying  bottle  b  containing 
concentrated  H2S04. 


CHLORINE 

Data  Necessary. 

Volume  of  flask,  V 

Volume,  corrected  to  o°  and  760  mm.,  VQ 

Weight  of  flask  filled  with  air,  W 

Weight  of  flask  filled  with  chlorine,  W 

Weight  of  vacuous  flask,  F 

Observed  temperature,  T 

Observed  pressure,  P 

Standard  temperature,  o°  C.,  T' 

Standard  pressure,  760  mm.,  P' 
Weight  of  a  liter  of  air  at  o°  and  760  mm.,  1.293  gms. 


39 


(X 


FIG.  12. 

Procedure.  Carefully  weigh  the  clean,  perfectly  dry 
flask  and  cork  (W) .  Generate  chlorine  in  the  apparatus 
shown  in  Fig.  12,  by  means  of  MnO2  and  concentrated 
HC1.  Let  the  chlorine  gas  run  into  the  previously 
weighed  flask  until  all  air  is  displaced,  i.e.,  until  the 


40  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

flask  is  entirely  filled  with  the  gas.  Remove  the  flask 
and  stopper  tightly.  Weigh  accurately  (W).  In  order 
to  ascertain  the  volume  of  the  flask  (F),  fill  it  with 
water  to  the  point  where  the  cork  comes  when  the  flask 
is  stoppered.  Then  carefully  measure  the  water  by 
means  of  a  graduated  cylinder.  This  gives  volume  of 
the  flask  (F).  Note  the  temperature  (T)  and  the 
pressure  (P)  in  the  room  at  the  time  the  experiment  is 
performed. 

Weight  of  the  Vacuous  Flask.  In  order  to  ascertain 
the  weight  of  the  chlorine,  it  is  necessary  to  know  the 
weight  of  the  vacuous  flask,  i.e.,  the  weight  when  it  is 
not  filled  with  air.  This  can  be  readily  calculated,  for 
we  know  the  weight  of  a  liter  of  air  at  o°  and  760  mm. 
pressure  to  be  1.293  gms.  and  consequently  the  weight 
of  i  cc.  of  air  under  these  conditions  is  0.001293  gm. 
The  capacity  of  the  flask  is  F  at  T  and  P.  Then  the 
capacity  at  o°  and  760  mm.  pressure  will  be 

VT'P 


Fn  = 


TP' 


This  corrected  volume,  Fo,  multiplied  by  0.001293  §m- 
gives   the  weight  of   the  air  in   the  flask,   and   W  • 
(Fo  X  0.001293)  equals  the  weight  of  the  vacuous  flask, 
or  F. 

Weight  of  a  Liter  of  Chlorine.  The  corrected  volume 
of  chlorine  is,  of  course,  the  same  as  that  obtained  for  air; 
or,  in  other  words, 

T/        vrp 
TPf  ' 

The  actual  weight  of  the  chlorine  in  the  flask  is  equal 
to  W'  —  F.     Knowing  that  F0  cc.  of  chlorine  weigh 


CHLORINE  41 

W'  —  F  gms.,  the  weight  of  a  liter  of  chlorine  can  be 
readily  calculated  from  the  proportion: 

Fo :  (W  -  F)  : :  1000  :  x, 

in  which  x  equals  the  weight  of  a  liter  of  chlorine. 

What  is  the  vapor  density  of  chlorine?  From  the 
vapor  density  calculate  the  theoretical  weight  of  a  liter 
of  chlorine  and  compare  the  result  with  that  obtained 
experimentally.  What  is  the  percentage  error? 

What  is  the  function  of  the  wash  bottle  a  of  the  chlorine 
apparatus  ?  What  is  the  function  of  bottle  b  ?  Mention 
two  possible  sources  of  error  in  the  determination  of 
the  weight  of  a  liter  of  chlorine  if  the  gas  used  was 
not  passed  through  the  wash  bottle  and  the  drying 
bottle. 

Hydrochloric  Acid  (HC1). 

64.  For  the  generation  of  hydrogen  chloride  (HC1), 
arrange  an  apparatus  as  shown  in  Fig.    n,  page  36. 
Place  about  20  gms.  of  NaCl  in  the  flask  and  add  through 
the  thistle  tube  100  cc.  of  a  mixture  of  equal  parts  by 
volume  of  water  and  concentrated  H2SO4.     (Pour  the 
acid  into  the  water.)     Warm  the  flask  gently  and  collect 
the  gas  evolved  by  displacement  of  air,  i.e.,  by  the 
same  method  used  for  the  collection  of  chlorine.     Collect 
four  or  five  bottles. 

Why  not  collect  the  gas  over  water? 

65.  Into  one  of  the  bottles  of  gas  thrust  a  burning 
splinter  to  ascertain  whether  the  gas  (HC1)  burns,  or 
supports  combustion.     What  reason  can  you  give  for 
the  action  ? 

66.  What  do  you  notice  when  a  bottle  of  the  gas  is 


42  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

uncovered  ?    Blow  across  the  open  mouth  of  a  bottle  of 
the  gas. 

Invert  a  bottle  of  the  gas  over  water  in  the  pneumatic 
trough.  Then,  letting  the  mouth  of  the  bottle  dip 
below  the  surface  of  the  water,  remove  the  cover  glass. 
What  causes  the  water  to  rise? 

67.  Pour  a  few  cubic  centimeters  of  ammonium  hy- 
droxide (NH4OH)  into  a  clean  bottle  and  shake  so  that 
the  inside  of  the  bottle  will  be  covered  with  the  liquid. 
Invert  the  bottle  over  one  of  the  bottles  containing  HC1 
gas  and  quickly  withdraw  the  cover  glass,  thus  bringing 
together  the  open  mouths  of  the  two  bottles. 

68.  Make  a  solution  of  HC1  gas  in  water  by  passing 
the  gas  into  about  100  cc.  of  cold  water  in  a  small  flask. 
The  delivery  tube  should  not  extend  to  the  bottom  of  the 
flask  but  should  barely  touch  the  surface  of  the  water. 
Why? 

In  separate  test  tubes  try  the  action  of  this  solution 
on  small  pieces  of  magnesium  (Mg),  zinc  (Zn),  iron  (Fe) 
and  sodium  carbonate  (Na^COs). 

Repeat  the  tests,  using  dilute  HC1  from  the  reagent 
shelf  instead  of  the  solution  made  above.  How  do  the 
results  compare? 

69.  Pour  a  few  cubic  centimeters  of  concentrated  H2S04 
into  a  test  tube  containing  about  5  cc.  of  concentrated 
HC1.    What  gas  is  evolved?    Why  is  it  evolved? 

70.  Into  separate  test  tubes  introduce -a  few  cubic 
centimeters  of  each  of  the  following  solutions:    silver 
nitrate   (AgNO3),  mercurous  nitrate   (HgNO3),  copper 
sulphate  (CuSO4),   lead  nitrate   (Pb(N03)2)   and  mag- 
nesium sulphate  (MgSO4).    To  each  tube  add  a  few 
cubic  centimeters  of  dilute  HC1. 


CHLORINE  43 

Which  metals  form  insoluble  chlorides? 

Summary.  What  is  the  odor  of  HC1?  When  we 
speak  of  " hydrochloric  acid"  as  a  laboratory  reagent, 
just  what  is  meant?  What  special  name  can  we  apply 
to  the  gas  to  distinguish  it  from  the  solution  ?  Speaking 
in  general  terms,  what  is  the  method  of  preparation  of 
HClgas? 

Problems,  (a)  Calculate  the  weight  of  7580  liters  of  HG1  gas 
at  10°  and  745  mm.  pressure.  If  this  amount  of  gas  is  dissolved 
in  the  proper  amount  of  water,  what  volume  of  15%  HC1  will  be 
formed,  the  specific  gravity  of  15%  HC1  being  1.075  ? 

(b)  How  many  liters  of  33%  HC1  (sp.  gr.  =  1.168)  can  be  pre- 
pared by  the  action  of  an  excess  of  H2S04  on  180  kilograms  of 
pure  sodium  chloride? 

Oxygen  Acids  of  Chlorine. 

71.  Hypochlorites.     Pass  chlorine  through  50  cc.  of 
a  solution  of  potassium  hydroxide  (KOH)  contained  in 
a  small  flask  until  the  liquid  is  saturated  with  the  gas. 
What  compound  is  formed  ? 

Introduce  a  portion  of  the  solution  thus  obtained 
into  a  small  beaker  and  immerse  a  piece  of  colored  calico 
in  it.  Allow  to  stand  for  a  few  minutes. 

72.  To  another  portion  of  the  solution  prepared  above, 
add  a  few  cubic  centimeters  of  dilute  H2S04.     What  gas 
is  evolved  ?    Repeat,  using  HC1  instead  of  H2SO4. 

73.  From  the  bottle  on  the  side  shelf  obtain  a  few 
grams    of    calcium    hypochlorite    (CaCl2O),    which    is 
commonly    called    " bleaching    lime"    or    "chloride    of 
lime."    Put  the  powder  in  a  small  beaker  and  add 
dilute  H2SO4. 

74.  In  a  bottle  or  tall  glass  cylinder  mix  a  few  grams 


44  EXPERIMENTS   IN  GENERAL  CHEMISTRY 

of  " bleaching  lime"  with  water.  Treat  the  mixture 
with  10  or  15  cc.  of  a  solution  of  cobalt  chloride  (CoCy 
and  allow  to  stand  10  minutes.  Notice  the  gas  evolved. 
Test  the  gas  with  a  glowing  splinter. 

75.  Chlorates.     Pass  chlorine  into  50  cc.  of  a  hot 
solution  of  KOH  until  the  liquid  is  saturated.     (Does 
chlorine  have  the  same  action  upon  hot  as  upon  cold 
KOH?)      Evaporate    the    solution    to    about    half    of 
its  original  volume  and  allow  to  stand  and  crystallize. 
If  crystals  do  not  separate  upon  cooling,  evaporate  the 
solution  to  a  still  smaller  volume  and  again  allow  to 
stand  and  cool,  quietly. 

What  is  the  composition  of  the  crystals  formed? 
Dry  the  crystals  by  pressing  them  between  pieces  of 
filter  paper,  and  reserve  them  for  later  experiments. 

76.  Into  a  dry  test  tube  introduce  a  very  small  crys- 
tal of  potassium  chlorate  and  treat  it  with  a  few  drops 
of  concentrated  H2S04.     (CAUTION.)     If  the  reaction  is 
not  noticeable  heat  the  test  tube  under  the  hood  being 
careful  to  protect  the  face  and  clothes. 

77.  Obtain  a  few  cubic  centimeters  of  pure  KC1O3  solu- 
tion from  the  bottle  on  the  side  shelf.    To  this  solution 
add  a  few  drops  of  a  solution  of  silver  nitrate  (AgNO3). 
Also  test  a  few  cubic  centimeters  of  potassium  chloride 
(KC1)  solution  with  AgNO3  solution.     Is  the  chlorine 
in  KC1O3  in  a  different  state  of  combination  from  that 
in  KC1? 

78.  Into  a  test  tube  partially  filled  with  pure  KC103 
solution  introduce  a  piece  of  zinc  and  enough  concen- 
trated H2SO4  to  cause  an  evolution  of  hydrogen.    Allow 
to  stand  for  5  minutes;  then  filter  and  test  the  clear  ni- 
trate with  AgN03  solution.     Compare  with  the  previous 


CHLORINE  45 

experiment.      What  change  has   the  zinc  and  H2SO4 
brought  about  ? 

79.  Perchlorates.     To  a  small  crystal  of  potassium 
perchlorate  (KC104)  in  a  test  tube  add  a  few  drops  of 
concentrated  H2S04.     Compare  with  Experiment  76. 

In  another  test  tube  try  the  action  of  concentrated 
HC1  on  a  few  crystals  of  KC104.  Is  chlorine  evolved  ? 
What  is  the  action  of  HC1  on  KC1O3  ? 

80.  Heat  a  few  crystals  of  KC1O4  in  a  hard  glass  test 
tube  and  test  for  evolved  oxygen  by  means  of  a  glowing 
splinter. 

81.  Immerse  a  strip  of  colored  calico  in  a  solution  of 
KC104.     Is  the  cloth  bleached  ?     Repeat  with  a  solution 
of  KClOs. 

82.  Test  the  action  of  a  solution  of  AgNOs  on  a  solu- 
tion  of  KC104   and   compare   the   results  with   those 
obtained  in  Exp.  77.     Repeat  Exp.  78,  using  a  solution 
of  KC104  instead  of  KC103. 

Summary.  What  is  the  relation  between  "  bleach- 
ing lime"  and  sodium  hypochlorite ?  How  do  these 
substances  bleach?  What  element  have  we  studied 
which  bleaches  in  the  same  way  ?  What  is  the  composi- 
tion of  "eau  de  Javelle"? 

Mention  ways  in  which  hypochlorites,  chlorates  and 
perchlorates  differ  in  their  action  with  reagents.  Why  is  it 
that  solutions  of  chlorates  and  perchlorates  do  not  bleach? 

Compare  the  action  of  concentrated  HC1  and  con- 
centrated H2S04  on  hypochlorites,  chlorates  and  per- 
chlorates and  write  general  equations-  for  each,  letting 
M  represent  a  monovalent  metal. 

Which  oxygen  acid  of  chlorine  have  we  not  studied? 
Why  has  it  been  omitted  ? 


46  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

BROMINE  (Br;  80). 

83.  Introduce  a  drop  of  bromine  into  a  clean,  dry 
500-cc.  flask  and  warm  gently  by  holding  at  some  dis- 
tance above  the  flame.     What  is  the  color  of  the  vapor 
of  bromine?   Has  bromine  any  odor?    (CAUTION!   Great 
care  must  be  used  in  handling  bromine;  it  is  a  dangerous 
chemical  and  its  vapors  are  very  irritating.     Do  not  get 
bromine  on  the  hands.) 

84.  Introduce   a  mixture   of   about   i    gm.   each   of 
Mn(>2  and  potassium  bromide  (KBr)  into  a  test  tube 
and   add   a   few   cubic   centimeters   of   strong   H2S04. 
Warm  the  mixture  gently  and  notice  the  gas  evolved. 
How  can  you  describe  it?     What  other  element  have 
you  prepared  by  heating  one  of  its  compounds  with 
Mn02  and  concentrated  H2SO4? 

85.  To  a  few  cubic  centimeters  of  a  solution  of  KBr 
in  a  test  tube  add  chlorine  water.     Does  the  liquid 
change  in  appearance?    Why?    Which  element,  Cl  or 
Br,  has  the  stronger  affinity  for  potassium  (K)  ? 

86.  Into  each  of  three  test  tubes  introduce  a  few  cubic 
centimeters  of  bromine  water  from  the  bottle  on  the 
side   shelf.     To   one  add  a  few  cubic   centimeters   of 
carbon  disulphide  (082) ;  to  the  second,  a  few  cubic  centi- 
meters of  chloroform  (CHC13);   and  to  the  third  a  few 
cubic  centimeters  of  ether  (C4Hi0O).     Shake  each  tube 
gently  and  then  allow  to  stand  for  a  few  minutes.     Notice 
what  has   taken  place.     Is   bromine   more   soluble  in 
water  or  in  the  reagents  used  ? 

87.  In  a  beaker  immerse  a  strip  of  colored  calico  in 
some  bromine  water.     Allow  to  stand  for  a  few  moments. 

In   separate  test  tubes  try  the   action   of  bromine 


BROMINE  47 

water   on   solutions   of   indigo,    cochineal   and    litmus. 
Compare  the  results  with  those  obtained  in  Exp.  62. 

88.  Under  a  hood  having  a  good  draught  pour  a  few 
cubic  centimeters  of  Br  into  a  small  beaker  or  a  wide- 
mouth  bottle.     Drop  a  small  piece  of  tin  (Sn)  foil  into 
the  beaker.     Explain  all  phenomena.     What  compound 
is  formed? 

Repeat,  using  a  red-hot  piece  of  thin  copper  foil  instead 
of  tin  foil. 

89.  Into  a  solution  of  NaOH  pour  a  few  cubic  centi- 
meters of  Br  water  and  stir.     Is  the  solution  brown? 
What  change  has  taken  place?     What  would  have  been 
formed  if  a  hot  solution  of  NaOH  had  been  employed? 
(Compare  with   experiments   on   the   oxygen   acids   of 
chlorine.) 

Put  a  piece  of  colored  calico  into  the  solution  formed 
by  adding  Br  water  to  NaOH.  Allow  to  stand  for  a 
few  minutes.  Is  the  color  changed? 

Hydrobromic  Acid  (HBr). 

90.  To  a  few  crystals  of  KBr  in  a  small  flask  add  a 
little  concentrated  H2S04  and  warm.       Breathe  across 
the  mouth  of  the  flask.     How  many  different  products 
of  the  reaction  can  you  identify?     What  are  the  brown 
fumes  which  finally  appear  in  the  flask?     Why  is  this 
not  a  good  method  for  the  preparation  of  HBr  ? 

91.  Arrange  an  apparatus  as  shown  in  Fig.  13,  using 
a  250  cc.  flask  and  a  small  beaker  of  water.      The  stem 
of  the  retort  should  barely  dip  beneath  the  surface  of  the 
water.     Be  sure  that  all  connections  are  tight.     Into 
the  flask  introduce  5  gms.  of  red  phosphorus  (P)  and 
20  cc.  of  water.     The  "U"  tube  should  be  filled  with 


48  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

pieces  of  pumice  coated  with  a  mixture  of  red  phospho- 
rus and  water. 

Through  the  dropping  funnel  now  introduce  15  cc.  of 
bromine,  adding  it  a  drop  at  a  time  and  agitating  the 
flask  after  each  addition.  The  HBr  generated  dissolves 
in  the  water  in  the  beaker  forming  a  solution  of  hydro- 
bromic  acid. 


FIG.  13. 

Test  the  solution  with  litmus  paper.  Try  its  action 
on  Na^COs  and  on  zinc  or  magnesium.  Is  it  an  acid  ? 

92.  In  separate  test  tubes  try  the  action  of  HBr 
upon  solutions  of  silver  nitrate  (AgNO3);  lead  nitrate 
(Pb(NO3)2  and  mercurous  nitrate  (HgN03). 

Repeat  the  experiment,  using  a  solution  of  KBr  instead 
of  HBr.  Does  it  make  any  difference  which  is  used  ? 

How  do  the  insoluble  bromides  compare  with  the  in- 
soluble chlorides? 


IODINE  49 

IODINE  (I;  127). 

93.  Introduce  a  small  crystal  of  iodine  into  a  clean, 
dry  5<x>-cc.  flask  and  warm  gently.     Note  the  color  of 
the  iodine  vapor.    When  all  the  iodine  is  vaporized, 
allow  the   flask  to  stand   and  cool.     Note  the   black 
deposit  on  the  sides  of  the  flask. 

Compare  the  color  of  iodine  and  its  vapor  with  bro- 
mine and  chlorine. 

94.  Make  a  mixture  of  a  little  MnO>  and  potassium 
iodide  (KI).    Place  the  mixture  in  a  test  tube,  add  a 
little  concentrated  H2SC>4  and  warm  gently.    What  is 
the  result? 

95.  In  separate  test  tubes  add  a  little  iodine  solution 
to   each   of    the   following:    CSa,    CHCls,   alcohol   and 
ether.     Shake  each  tube  gently  and  then  allow  to  stand. 
What  colors  are  produced?     Compare  with  Exp.  86. 
Which  solvents  dissolve  iodine  to  give  a  solution  the 
color  of  iodine  vapor? 

96.  To  a  few  cubic  centimeters  of  KI  solution  add  a 
little  chlorine  water;   then  add  a  few  cubic  centimeters 
of  CS2  and  shake. 

Repeat,  using  bromine  water  instead  of  chlorine. 
What  do  you  conclude  as  to  the  relative  affinity  of 
chlorine,  bromine  and  iodine  for  potassium? 

97.  To  about  2  cc.  of  KI  solution  add  strong  chlor- 
ine water  until  in  decided  excess.    What  great  change 
has  taken  place  ?     Explain  fully. 

98.  In  a  test  tube  treat  a  small  crystal  of  iodine  with 
alcohol.    Note    that    the    iodine    dissolves.    What    is 
"  tincture  of  iodine "?      (Save  the  solution  for  use  in 
Exp.  101.) 


50  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

99.  Try  the  action  of  iodine  water  on  starch  paste. 
Have  you  seen  this  color  in  any  previous  experiments? 
Heat  the  solution  to  boiling  and  then  allow  to  stand  and 
cool.    Note  all  changes. 

100.  In  a  test  tube  treat  a  crystal  of  iodine  with  water 
and    shake.     Does     the    iodine    dissolve?     To    what 
extent  ? 

Now  add  a  crystal  of  KI  and  again  shake.  Does  the 
addition  of  KI  cause  any  hastening  of  solution  ?  Why  ? 

101.  To  a  few  cubic  centimeters  of  tincture  of  iodine 
add  a  few  drops  of  water.     Why  is  there  a  change  ? 

102.  To  a  beaker  containing  about  50  cc.  of  water  add 
enough  of  the  solution  formed  in  Exp.  100  to  produce 
a  faint  yellow  color.     Add  a  drop  or  two  of  starch 
paste.     What  happens  ?     Now  add  a  solution  of  sodium 
thiosulphate  (Na^Os),  a  drop  at  a  time,  until  in  excess. 
Explain  all  changes. 

Hydriodic  Acid  (HI). 

103.  To  a  few  crystals  of  KI  in  a  small  flask  add  con- 
centrated H2S04  and  warm  gently.     Breathe  over  the 
mouth  of  the  flask;   are  there  any  white  fumes?     Con- 
tinue to  heat  gently  and  notice  all  the  products  formed 
by  the  reaction.     Explain  by  equations. 

104.  To  about  50  cc.  of  water  in  a  small  beaker  or 
flask   add   a   little   powdered   iodine.     Pass   hydrogen 
sulphide  (H2S)  gas  through  the  solution  until  all  iodine 
disappears.     What    is    the    white    substance    formed? 
What  is  in  solution? 

Filter  the  solution  and  test  the  clear  nitrate  with 
blue  litmus  paper.  In  separate  test  tubes  try  the  action 
of  the  solution  on  zinc  and  on 


FLUORINE   AND    HYDROFLUORIC   ACID  5 1 

105.  In  separate  test  tubes  add  KI  solution  to  solu- 
tions of  AgNO3,  HgNO3,  Hg(N03)2,  NiS04  and  Pb(NO3)2. 
Which  metals  form  insoluble  iodides  ? 

FLUORINE  (F;    19). 

Hydrofluoric   Acid    (HF). 

1 06.  Coat  the  concave  side  of  a  watch  glass  with  par- 
affin by  warming  gently  and  rubbing  with  a  small  piece 
of  paraffin.     When  cold,  scratch  a  design  in  the  paraffin. 
Carefully  pour  a  few  cubic  centimeters  of  hydrofluoric 
acid  (HF)  on  the  glass  and  hold  it  so  that  the  acid  will 
come  into  contact  with  the  glass  where  the  paraffin  has 
been  scratched  away. 

Remove  the  excess  of  paraffin,  after  washing  away  the 
acid,  by  warming  the  glass  or  by  wiping  it  off  with 
a  cloth  wet  with  alcohol.  Examine  the  clean  glass. 
(CAUTION!  HF  is  a  very  dangerous  chemical.  Do  not 
get  it  on  the  hands  —  it  makes  very  bad  sores.  Do  not 
breathe  the  vapors  —  they  are  very  poisonous?) 

107.  Coat  one  side  of  a  piece  of  window  glass  with 
paraffin  and  scratch  a  design  through  the  latter. 

Place  about  2  gms.  of  calcium  fluoride  (CaF2),  com- 
monly called  "fluorite"  or  "  fluorspar, "  in  a  lead  dish 
and  moisten  with  enough  concentrated  H2S04  to  form 
a  thick  paste.  Place  the  glass,  paraffin  side  down,  over 
the  dish  and  allow  to  stand  overnight. 

Then  clean  the  glass  and  notice  the  design  etched 
upon  the  surface. 

1 08.  Into  a  test  tube  introduce  a  mixture  of  about 
i  gm.  each  of  sand  or  powdered  quartz  (Si02)  and  CaF2. 
Add  a  little  strong  H2S04.    Warm  and  at  the  same  time 
hold  a  glass  rod  with  a  drop  of  water  at  the  end,  in  the 


52  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

mouth  of  the  test  tube.  What  causes  the  water  to  be- 
come turbid?  For  what  is  this  a  good  qualitative  test? 

109.  In  separate  test  tubes  try  the  action  of  sodium 
fluoride  (NaF)  solution  on  solutions  of  AgN03  and 
Ca(OH)2. 

For  the  sake  of  comparison,  try  the  action  of  NaCl 
on  solutions  of  AgN03  and  Ca(OH)2. 

Summary.  To  what  family  of  elements  does  fluorine 
belong?  What  other  elements  are  included  in  this 
family?  Write  a  general  equation  for  the  preparation 
of  the  halogens,  letting  X  represent  halogen.  Can  flu- 
orine be  prepared  by  heating  one  of  its  salts  with  MnC>2 
and  H2S04?  Why?  How  is  it  possible  to  prepare  the 
element  fluorine? 

In  what  respects  does  fluorine  differ  from  the  other 
halogens  ?  Why  is  hydrofluoric  acid  always  kept  in  wax 
or  lead  bottles  ? 

Problems,  (a)  The  specific  gravity  of  sea  water  is  1.025,  and  it 
contains  0.36  part  of  MgBr2  per  1000.  How  many  cubic  centi- 
meters of  bromine  can  be  obtained  from  i  cubic  meter  of  sea 
water.  Specific  gravity  of  Br  =  3.18. 

(b)  To  make  12  liters  of  a  35%  solution  of  sodium  hypobromite 
(specific  gravity  =  1.24),  how  many  grams  of  NaOH  and  how 
many  cubic  centimeters  of  bromine  are  necessary? 

(c)  How  many  grams  of  iodine  can  be  obtained  as  a  by-product 
from  10  tons  of  "caliche"  (Chili  saltpeter)  containing  1.3%  of 
sodium  iodate? 

(d)  1 8  gms.  of  iodine  will  occupy  what  volume  if  vaporized  at 
40°  and  765  mm.  pressure? 

(e)  How  many  grams  of  17%  HF  solution  can  be  prepared  from 
325  gms.  of  fluorspar? 


CHAPTER  VI. 

ACIDS,   BASES   AND    SALTS. 

no.  To  a  beaker  of  water  add  a  few  drops  of  concen- 
trated HCL  Taste  the  solution  thus  formed.  Try  the 
action  of  the  solution  on  litmus  paper,  on  turmeric  paper 
and  on  a  solution  of  phenolphthalein. 

in.  Repeat  the  preceding  experiment,  using  a  few 
drops  of  NaOH,  a  base,  instead  of  HCL  Test  as  before. 

112.  In  a  porcelain  evaporating  dish  add  dilute  HC1 
to  20  cc.  of  dilute  NaOH  until  the  solution  is  neutral  to 
litmus  paper.    Take  out  the  litmus  paper  and  then 
evaporate  the  solution  to  dryness.    Taste  the  residue. 
What  sort  of  a  compound  is  it? 

113.  By  means  of  a  deflagration  spoon  burn  a  small 
piece  of  metallic  sodium  (Na)  in  a  5oo-cc.  flask.     Add  a 
little  water,  shake  the  flask  and  then  test  the  reaction 
of  the  water  towards  litmus.     Does  the  water  contain 
an  acid  or  a  base  ? 

Repeat  the  experiment,  burning  a  small  piece  of  phos- 
phorus (a  non-metal)  instead  of  the  sodium.  Is  a  base 
or  an  acid  formed  in  this  experiment  ?  (The  flask  should 
be  thoroughly  cleaned  before  this  second  part  of  the 
experiment  is  performed.) 

The  oxide  of  a  metal  +  water  forms  what  kind  of  a 
compound  ?  The  oxide  of  a  non-metal  +  water  forms 
what  kind  of  a  compound  ? 

114.  Boil  a  little  NaOH  solution  and  test  the  vapors 
with  wet  turmeric  paper.     Is  there  any  change  in  color  ? 
Repeat,  using  NH4OH  instead  of  NaOH. 

53 


54  EXPERIMENTS  IN  GENERAL  CHEMISTRY 

How  does  ammonium  hydroxide  differ  from  other 
bases  in  its  physical  properties? 

Acid,  Basic  and  Neutral  Salts. 
BASICITY  OF  AN  ACID. 

115.  Pour  a  little  SbCl3  solution  into  a  small  beaker 
of  water.     What  happens?    What  kind  of  a  compound 
is  the  white  precipitate?    What  happens  when  concen- 
trated HC1  is  added  to  the  white  precipitate  ? 

1 1 6.  In  a  porcelain  evaporating  dish  carefully  neutral- 
ize 20  cc.  of  dilute  H2SC>4  with  NaOH.     Evaporate  to 
dryness  and  heat  gently.    Dissolve  the  residue  in  a  little 
water  and  test  the  action  of  the  solution  towards  litmus. 

Now  add  another  20  cc.  of  dilute  H2S04  to  the  solution 
and  again  evaporate  to  dryness.  Dissolve  the  residue 
in  water  and  test  its  reaction  towards  litmus. 

Do  the  two  parts  of  this  experiment  yield  different 
results?  How,  can  you  account  for  it?  What  kind  of  a 
salt  is  the  one  obtained  last?  What  is  the  basicity  of 
H2S04? 

Repeat,  using  HNOs  in  each  case  instead  of  H2SO4. 
Are  two  different  products  formed?  What  is  the  ba- 
sicity of  HNO3? 

117.  Test  the  reaction  towards  litmus  paper  of  solu- 
tions of  the  following  compounds:    CuS04,    Na2HPO4 
NaHCO3,  Na2CO3,  ZnSO4,  NaCl  and  MgS04. 

Do  all  neutral  salts  have  a  neutral  reaction?  Which 
of  the  above  are  exceptions? 

Do  all  acid  salts  have  an  acid  reaction?  Which  are 
exceptions  ? 

Name  a  neutral  salt  that  has  an  acid  reaction,  one  that 
has  a  neutral  reaction  and  one  that  has  a  basic  reaction. 


ACIDS,   BASES   AND   SALTS  55 

Summary.  Do  all  acids  contain  oxygen  ?  What  one 
element  is  always  present  in  an  acid?  What  name  is 
applied  to  those  acids  which  contain  oxygen  ? 

What  is  characteristic  about  the  formula  for  an  acid  ? 
For  a  base  ?  Is  water  an  acid  or  a  base  ?  Why  ?  What 
is  characteristic  about  the  formula  for  an  acid  salt? 
For  a  basic  salt  ? 

Which  of  the  following  are  acids,  which  are  bases  and 
which  are  salts:  H3PO4,  Nal,  HI,  H2MoO4,  A1(OH)8, 
H2S04,  Fe(OH)3,  FeSO4,  H3BO3,  NH4OH? 

Which  of  the  following  do  you  consider  as  acid  salts, 
which  are  neutral  salts  and  which  are  basic  salts:  Na3P04, 
NaHCO3,  Pb(OH)(N03),  CaH2(CO3)2,  Na2HPO4,  NaCl, 
CaC03,  NaHS04,  NaKS04,  Mn2(OH)2C03? 

What  is  meant  by  the  basicity  of  an  acid?  Give 
examples  of  mono-,  di-,  and  tribasic  acids.  What  is 
meant  by  the  acidity  of  a  base  ? 

ACIDIMETRY  AND  ALKALIMETRY. 
(Quantitative.) 

1 1 8.  Obtain  from  the  instructor  a  supply  of  standard 
H2SO4  solution  of  known  strength.  What  weight  of 
absolute  H2SO4  does  each  cubic  centimeter  of  this  solu- 
tion contain?  To  what  weight  of  NaOH  is  each  cubic 
centimeter  equivalent? 

Rinse  out  a  burette  twice  with  small  portions  of  the 
acid  solution  and  then  completely  fill  the  burette  with 
the  same  solution.  In  a  clean  flask  obtain  from  the 
instructor  an  unknown  sample  of  NaOH  solution  to  test 
by  means  of  the  acid  solution  in  the  burette.  This 
method  of  testing  is  known  as  "  titration  "  and  is  per- 
formed as  follows:  Add  to  the  contents  of  the  flask 


EXPERIMENTS  IN   GENERAL  CHEMISTRY 


two  or  three  drops  of  phenolphthalein  solution  to  serve 
as  an  "indicator."  The  alkali  turns  the  phenolphthalein 
bright  red. 

Read  the  level  of  the  acid  in  the 
burette  and  then  carefully  allow 
the  acid  to  run  drop  by  drop 
into  the  flask  containing  the  alkali 
and  indicator,  until  the  alkali  is 
neutralized.  When  the  alkali  is 
completely  neutralized,  the  addi- 
tion of  one  more  drop  of  acid  will 
completely  discharge  the  red  color; 
therefore  continue  to  add  acid  un- 
til one  drop  finally  causes  the  red 
color  to  disappear.  Again  read 
the  level  of  the  acid  in  the  burette 
and  calculate  the  weight  of  NaOH 
which  was  in  the  solution  in  the 
flask.  (Fig.  14.) 

Obtain  two  other  samples  of  un- 
known alkali  from  the  instructor 
and  titrate  in  the  same  manner. 

Standardization  of  an  Alkali 
Solution.  Obtain  from  the  in- 
structor a  supply  of  the  alkali 
solution  to  be  standardized.  Fill 
the  other  burette  with  this  solution 
after  having  first  rinsed  it  out  twice 
with  small  portions.  Into  a  clean  flask  carefully  meas- 
ure out  from  the  burette  10  cc.  of  the  alkali  solution. 
Add  two  or  three  drops  of  the  indicator  and  then  titrate 
with  the  known  acid  solution  in  the  other  burette  as 


FIG.  14. 


ACIDS,    BASES   AND   SALTS  57 

described  previously.  Make  three  titrations,  take  the 
average  reading  and  calculate  the  strength  of  the  alkali 
solution  and  its  H2S04  equivalent  per  cubic  centimeter. 

Now  obtain  from  the  instructor  samples  of  unknown 
acid  to  determine  by  means  of  this  "standardized" 
alkali  solution.  The  method  of  titration  is  practically 
the  same,  though  the  solution  after  the  addition  of  the 
indicator  is  colorless  and  is  titrated  with  the  alkali  until 
a  drop  of  the  latter  finally  produces  a  faint  permanent 
pink  color.  The  point  at  which  the  indicator  shows  a 
change  in  color  is  called  the  "end  point." 

For  further  exercises  in  "acidimetry  and  alkalimetry" 
obtain  instructions  from  the  instructor. 


CHAPTER  VII. 
NITROGEN  (N;  14). 

119.    Preparation.     Place   a   small    piece   of    phos- 
phorus on  a  porcelain  crucible  cover  or  on  a  piece  of 


FIG.  15. 

cork  floating  on  water.  Ignite  the  phosphorus  by  touch- 
ing it  with  a  hot  file,  and  quickly  cover  with  an  inverted 
beaker  or  bottle,  allowing  the  latter  to  dip  into  the 
water. 

58 


NITROGEN 


59 


Note  the  white  fumes.  When  combustion  is  com- 
plete, note  the  decrease  in  volume  of  the  gas  in  the 
bottle.  What  does  this  decrease  represent?  Allow  the 
bottle  to  stand  in  the  water  for  a  time.  Do  the  white 
fumes  disappear?  Why? 

Test  the  gas  remaining  in  the  bottle  as  directed  in 
Exp*  121. 

120.  Arrange  an  apparatus  as  shown  in  Fig.  15.     Into 
the  flask  put  25  cc.  of  water,  10  gms.  sodium  nitrite 
(NaNC>2)   and    5    gms.   ammonium    chloride   (NH4C1). 
Heat  the  mixture  gently  and  collect  two  or  three  bottles 
of  the  gas.     Test  the  gas  as  directed  in  the  following 
experiment. 

121.  Properties.    Test  the  samples  of  nitrogen  pre- 
pared in  the  two  preceding  experiments  by  introducing 
a  burning  splinter  into  the  bottles.     Does  nitrogen  burn  ? 
Does  it  support  combustion?     Do  the  two  samples  of 
nitrogen  give  the  same  test? 

Has  the  gas  any  odor?  Is  nitrogen 
soluble  in  water?  Has  it  any  reaction 
towards  litmus  paper? 

Determine  whether  nitrogen  will  sup- 
port the  combustion  of  sulphur. 


Air. 

122.   Tests  for  Impurities.     Arrange  a 
bottle  as  shown  in  Fig.  16.     Introduce  FIG.  16. 

lime  water  or  baryta  water  into  the 
bottle  and  draw  air  through  the  solution  by  attaching 
the  bottle  to  the  suction  pump.     Allow  this  experiment 
to  run  for  some  time.    Test  the  precipitate  formed  by 
adding  HC1.     If  this  produces  an  effervescence,  the  pre- 


60  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

cipitate  was  a  carbonate,  hence  the  solution  took  carbon 
dioxide  from  the  air.  What  do  you  know  of  the  relative 
amounts  of  this  impurity  in  the  air  in  the  country  and 
in  the  city? 

By  what  means  can  you  detect  moisture  in  the  air? 
What  are  some  of  the  other  impurities  always  present 
in  the  air?  Mention  several  other  impurities  which 
may  be  present  in  the  air  in  the  laboratory.  How  can 
you  detect  these? 

ANALYSIS  or  AIR. 
(Quantitative.) 

123.  Procuring  the  Sample.  Disconnect  the  appa- 
ratus (Fig.  17)  at  point/.  Open  the  stopcock  in  tube  a 
of  the  gas  burette  and  raise  tube  b  until  a  is  completely 
filled  with  water.  Now  lower  b,  thus  allowing  air  to  enter 
a.  Hold  the  two  tubes  in  such  a  way  that  the  water  is 
at  the  same  level  in  both  and  a  contains  just  50  cc.  of 
air.  While  still  holding  the  tubes  in  this  position,  close 
the  stopcock  in  a. 

Connecting  the  Apparatus.  Gently  blow  into  the 
rubber  tube  d  attached  to  the  gas  pipette,  thus  forcing 
the  alkaline  pyrogallic  acid  solution  contained  in  c  up 
through  the  capillary  tube  e.  When  e  is  completely 
filled  with  the  liquid,  pinch  the  rubber  tube  d  to  pre- 
vent the  liquid  from  flowing  back,  and,  while  doing  this, 
connect  e  to  a  at  point  /  by  means  of  a  short  piece  of 
rubber  tubing,  as  shown  in  Fig.  17.  The  pressure  on 
rubber  tube  d  may  then  be  released. 

If  everything  has  been  done  correctly  up  to  this  point, 
tube  a  will  contain  50  cc.  of  air,  and  bulb  c  and  tube  e 


NITROGEN 


61 


will  be  completely  filled  with  the  alkaline  pyrogallic 
acid  solution. 

Absorption  of  the  Oxygen.    Open  the  stopcock  in  a 
and  elevate  b  in  order  to  force  all  of  the  air  in  a  through 


FIG.  17. 

tube  e  into  the  gas  pipette.  As  soon  as  all  of  the  air  is 
in  c,  close  the  stopcock  in  a.  (Tube  a  and  tube  e  will 
then  be  completely  filled  with  water.) 

Without  disconnecting  the  apparatus,  gently  agitate 
the  pipette  in  order  to  bring  the  air  into  better  contact 
with  the  alkaline  pyrogallic  acid  solution.  Continue  to 
shake  for  about  10  minutes.  Then  open  the  stopcock  in 
a  and  lower  6,  thus  causing  the  air  to  be  driven  back  into 
the  tube  a.  As  soon  as  all  air  has  reentered  a,  and 
tube  e  is  completely  filled  with  the  solution  from  c,  close 


62  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

the  stopcock  in  a.  Hold  a  and  b  in  such  a  manner  that  the 
water  in  them  will  be  at  the  same  level;  then  read  the 
volume  of  air  in  a. 

Force  the  air  into  c  a  second  time  and  shake  for  about 
5  minutes,  after  which  drive  the  air  back  into  a  and  read 
the  volume  as  before.  This  process  should  be  continued 
until  two  successive  readings  are  identical.  In  reading 
the  volume  of  the  gas  in  a,  always  hold  the  tubes  in 
such  a  way  that  the  water  is  at  the  same  level  in  both. 
(Why?) 

Results.  By  the  treatment  described  above,  all  the 
oxygen  in  the  sample  of  air  is  dissolved  by  the  alkaline 
solution  of  pyrogallic  acid  and  the  nitrogen  alone  re- 
mains. From  your  readings  calculate  the  percentage 
by  volume  of  oxygen  and  nitrogen  in  the  air. 

Mention  several  gases  which  occur  in  the  air  in  minute 
quantities,  and  tell  which  affect  the  result  for  oxygen 
obtained  above  and  which  affect  the  nitrogen  percen- 
tage. 

Ammonia. 

124.  Test.     Hold  a  piece  of  moist  red  litmus  paper 
near  an  open  bottle  of  ammoniiim  hydroxide  and  notice 
the   change   in   color.     Repeat,   using   moist    turmeric 
paper  instead  of  litmus. 

Notice  the  odor  of  ammonium  hydroxide.  Why  has 
this  liquid  an  odor?  Why  are  the  litmus  and  turmeric 
paper  changed  when  brought  near  the  open  bottle  ? 

125.  Preparation.     For  the  preparation  of  ammonia 
use  the  apparatus  shown  in  Fig.  18.     The  drying  tube 
should  be  filled  with  small  dry  pieces  of  soda-lime.     Into 
the  flask  introduce  a  mixture  of  about  30  gms.  each  of 
NH4C1  and  slaked  lime,  Ca(OH)2.    Heat  gently  and 


NITROGEN 


collect  the  gas  by  displacement  of  air.  Collect  several 
bottles,  cover  with  glass  plates,  and  save  for  use  in 
Exp.  130. 


FIG.  18. 

126.  In  a  test  tube  gently  warm  a  mixture  of  NaOH  or 
KOH  and  a  solution  of  some  ammonium  salt.    Hold  a 
piece  of  moistened  litmus  paper  near  the  mouth  of  the 
test  tube.     Repeat  with  another  ammonium  solution. 

127.  In  separate  test  tubes  heat  bits  of  leather,  glue 
or  egg  albumen  with  small  amounts  of  soda-lime.     Is 
there  a  gas  evolved  ?    Note  the  odor  and  test  with  moist 
turmeric  paper. 


64  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

128.  In  a  hard  glass  test  tube  strongly  heat  about 
20  gms.  of  iron  turnings  or  filings  with  2  gms.  of  KN03. 
Test  the  escaping  gas  with  a  burning  splinter. 

In  like  manner  heat  20  gms.  of  iron  filings  with  2  gms. 
of  solid  KOH  and  test  the  escaping  gas  with  a  lighted 
splinter. 

Make  a  mixture  of  2  gms.  each  of  iron  turnings,  KN03 
and  KOH.  Heat  the  mixture  in  a  test  tube  and  test 
the  escaping  gas  with  turmeric  paper.  Has  the  gas  an 
odor? 

Explain  fully  the  results  obtained  from  the  three  parts 
of  this  experiment. 

129.  In  a  test  tube  dissolve  a  small  crystal  of  KNO3 
in  about  5  cc.  of  strong  KOH  solution.     Add  a  little 
metallic  aluminum  (wire  or  turnings)  and  heat.     Test 
the  escaping  gas  with  wet  turmeric  paper. 

130.  Properties.     Test  a  bottle  of  ammonia  gas  with 
a  burning  splinter.     Is  the  gas  combustible?     Does  it 
support  combustion?     Open  a  bottle  of  the  gas  under 
water.     Does  the  water  rise  in  the  bottle?    What  can 
you  say  as  to  the  solubility  of  NH3?    What  other  very 
soluble  gas  have  we  studied  ? 

Hydroxylamine. 

131.  In  separate  test  tubes  try  the  action  of  hydroxyl- 
amine  hydrochloride  solution  on  solutions  of  mercuric 
chloride  (HgCl2)  and  copper  sulphate  (CuSO4).     What 
kind  of  an  action  has  the  reagent  on  these  solutions? 

What  is  the  formula  for  hydroxylamine  hydrosulphate  ? 
What  other  compounds  of  nitrogen  and  hydrogen  have 
you  studied  ?  Tell  the  properties  of  each. 


NITROGEN  65 

DETERMINATION  or  THE  WEIGHT  OF  A  LITER 
OF  AMMONIA. 

(Quantitative.) 

132.  Proceed  exactly  as  described  in  the  experiment 
on  the  determination  of  the  weight  of  a  liter  of  chlorine 
(Exp.  63,  page  38).     The  ammonia  can  be  generated 
by  means  of  the  apparatus  as  described  in  Exp.  125, 
or  it  may  be  generated  by  simply  boiling  a  strong  solu- 
tion of  ammonium  hydroxide.     In  both  cases,  the  gas 
should  be  dried  by  passing  through  a  tube  containing 
soda-lime. 

Why  not  dry  the  gas  with  concentrated  sulphuric  acid 
as  in  the  case  with  chlorine  ?  Why  is  it  not  necessary 
to  wash  the  ammonia  as  was  done  with  chlorine  ? 
What  would  happen  if  this  were  done  ? 

Nitrous  Oxide  (N2O). 

133.  Arrange  an  apparatus  as  shown  in  Fig.  19,  using 
a  flask  of  about  250  cc.  capacity.     Introduce  about  25 
gms.  of  dry  ammonium  nitrate  (NI^NOs)  into  the  flask 
and  heat  gently.    After  the  gas  begins  to  be  evolved 
steadily  and  all  air  has  been  driven  from  the  apparatus, 
collect  several  bottles  of  the  gas  by  displacement  of  water. 

134.  Introduce  a  glowing  splinter  into  a  bottle  of  the 
gas.    What  other  gas  have  you  studied  which  affects  a 
glowing  splinter  in  the  same  way?     Mention  several 
tests  by  which  these  two  gases  can  be  distinguished. 

135-  By  means  of  a  deflagrating  spoon  introduce  into 
a  bottle  of  the  gas  a  little  sulphur  which  is  burning  only 
feebly.  Repeat  the  test,  first  heating  the  sulphur  highly 


66 


EXPERIMENTS  IN  GENERAL  CHEMISTRY 


so  that  it  burns  strongly.     Is  the  result  the  same? 
Why? 

136.  Burn  a  piece  of  phosphorus  in  the  gas.    What 
products  are  formed? 


X) 


FIG.  19. 

Invert  a  bottle  of  the  gas  in  water  and  allow  to  stand 
until  the  end  of  the  laboratory  period.  Does  the  water 
rise  in  the  bottle?  What  can  you  say  of  the  solubility 
of  nitrous  oxide  ? 


Nitric  Oxide  (NO). 

137.  Introduce  about  15  gms.  of  copper  turnings  into 
an  apparatus  as  shown  in  Fig.  2,  page  5.  Add  enough 
warm  water  to  cover  the  copper,  and  then  add,  through 
the  thistle  tube,  enough  concentrated  HNOs  to  cause  a 
brisk  evolution  of  gas.  When  all  air  has  been  driven 


NITROGEN 


67 


from  the  flask,  collect  several  bottles  of  the  gas  over 
water.    Note  the  color  of  the  gas  in  the  bottles. 

138.  Open  a  bottle  of  the  gas  so  that  it  comes  into 
contact  with  the  air.  What  pronounced  change  im- 
mediately takes  place? 


1 


FIG.  20. 


Test  the  gas  to  ascertain  if  it  will  support  combus- 
tion or  burn. 

139.  Test  with  burning  sulphur  as  described  in  Exp. 
135.  Does  it  make  a  difference  whether  the  sulphur  is 
burning  feebly  or  strongly  ? 

Likewise  test  with  phosphorus,  first,  burning  feebly, 
and  second,  burning  strongly.  Explain  all  phenomena 
observed. 


68  EXPERIMENTS  IN  GENERAL  CHEMISTRY 

140.  Pass  NO  into  a  solution  of  ferrous  sulphate 
(FeSO4)  in  a  test  tube.    When  the  solution  has  changed 
color,  remove  the  test  tube  and  heat  the  solution  to 
boiling.     Is  there  another  change? 

Nitrogen  Trioxide  (N203). 

141.  In  a  test  tube  gently  heat  a  mixture  of  a  few  grams 
of  arsenious  oxide  (A^Os)  and  a  few  cubic  centimeters 
of  concentrated  HNOs.     Notice  the  color  of  the  fumes 
produced.    If  the  fumes  were  condensed  to  a  liquid,  what 
would  be  the  color? 

142.  In  like  manner  heat  concentrated  HN03  with 
starch  in  a  small  flask  arranged  with  delivery  tube  as 
shown  in  Fig.  20.     Pass  the  gas  through  about  15  cc. 
of  water.     Test  the  reaction  of   this  aqueous  solution 
towards  litmus  paper.     Save  the  solution  for  use  in  Exp. 
144. 

Nitrous  Acid  (HN02). 

143.  Prepare  HNO2  by  passing  NO  through  10  cc.  of 
concentrated  HNOs  which  has  been  previously  diluted 
with  5  cc.  of  water.    Does  the  solution  change  in  color  ? 

144.  Try  the  action  of  the  nitrous  acid  thus  formed 
on  KMnO4  solution  and  on  KI  solution.     Add  a  little 
to  a  beaker  of  water  containing  a  few  drops  of  Kl-starch 
paste. 

Repeat  these  tests,  using  the  solution  prepared  in 
Exp.  142  by  passing  N2O3  into  water. 

145.  Sodium  Nitrite.     In  an  iron  dish  heat  a  mixture 
of  10  gms.  of  NaNO3  and  25  gms.  of  metallic  lead.     Allow 
the  resulting  dark  brown  mass  to  cool,  extract  with  water 
and  filter. 


NITROGEN  69 

Test  a  portion  of  the  solution  for  a  nitrite  by  adding 
a  little  dilute  H2S04  and  a  few  drops  of  Kl-starch  paste. 

To  another  portion  add  a  little  solid  NH4C1  and  warm. 
Test  the  gas  evolved.  What  is  it? 

Nitrogen  Tetroxide  (N2O4). 

146.  Heat   a   few   grams   of   powdered   lead   nitrate 
(Pb(NO3)2)  in  a  dry  test  tube  and  notice  the  colored 
fumes  produced.     Have  you  noticed  fumes  of  this  color 
before?     Mention  another  manner  in  which  this  gas, 
N2O4,  is  formed.     What  other  gas  has  the  same  color  as 
N204? 

Note  the  odor  of  N204,  but  breathe  very  little  of  the 
gas  inasmuch  as  it  is  poisonous. 

147.  Using  an  apparatus  as  shown  in  Fig.  20,  generate 
N204  by  heating  Pb(NOs)2  and  pass  the  gas  into  15  cc.  of 
water.     Does  the  water  change  in  color  ?     Does  the  gas 
appear  to  be  very  soluble  in  water  ? 

Test  the  aqueous  solution  thus  formed  by  means  of  lit- 
mus paper.  Is  it  an  acid  or  a  base  ?  What  compounds 
are  contained  in  the  water?  What  could  be  added  to 
prove  the  presence  of  one  of  these  ? 

Nitric  Acid  (H.NO3). 

148.  Introduce  25  gms.  of  NaNOs  and  15  cc.  of  con- 
centrated H2SO4  into  a  glass  stoppered  retort  arranged 
as  shown  in  Fig.  21,  with  a  test  tube  for  a  receiver.     The 
test  tube  is  cooled  in  a  dish  of  cold  water,  preferably  con- 
taining a  little  ice.     Heat  the  retort  and  collect  the  nitric 
acid  which  distills.     When  sufficient  acid  has  distilled, 
completely  remove  the  retort  and  test  tube  from  the  water. 


EXPERIMENTS   IN   GENERAL   CHEMISTRY 


149.  What  is  the  color  of  the  HN03  prepared  in  Exp. 
148?  To  what  is  the  color  due?  Using  a  long  glass 
tube,  bent  at  a  right  angle,  blow  through  the  acid  in  the 
test  tube  for  several  minutes.  Does  the  color  change  ? 

Drop  a  small  piece  of  copper  into  the  test  tube  of  acid. 
If  the  liquid  in  the  tube  is  HNO3,  what  will  be  produced 
when  it  comes  into  contact  with  the  copper  ? 


K) 


FIG.  21. 

150.  In  separate  test  tubes  try  the  action  of  HNOa 
on  zinc,  iron,  lead  and  tin.     Perform  these  tests  under 
the  hood.     Write  all  equations. 

151.  Treat  a  little  sulphur  with  concentrated  HNO3  in 
a  test  tube  and  boil  for  a  few  moments.     Does  the  sulphur 
dissolve?    What  has  become  of  it?    Add  a  few  drops 
of  BaCl2  solution  to  the  contents  of  the  test  tube  —  a 
white  precipitate  proves  the  presence  of  sulphuric  acid. 


NITROGEN  71 

152.  In  like  manner  boil  concentrated  HNOs  with  a 
little  phosphorus.     Test  a  portion  of  the  solution  formed 
with  ammonium  molybdate  solution.     This  reagent  gives 
a  yellow  precipitate  with  phosphoric  acid,  best  upon 
gently  warming. 

Explain  fully  the  manner  in  which  HNOs  oxidizes. 

153.  Place  a  drop  of  concentrated  HNOs  in  the  palm 
of  the  hand.     After  a  few  seconds  wash  off  the  acid  and 
treat  the  spot  with  NH4OH,  which  neutralizes  the  acid. 
Wash  and  dry  the  hand  and  examine  to  see  if  the  acid 
has  left  any  stain.    What  is  formed  by  the  action  of 
nitric  acid  on  flesh  ? 

What  is  guncotton  ?    Nitroglycerine  ? 

154.  Test  for  Nitric  Acid  and  Nitrates.    In  a  test 
tube  make  a  mixture  of  about  equal  parts  of  FeSO4  solu- 
tion and  a  solution  of  the  substance  to  be  tested  for 
nitrates.     Now  carefully  add  a  few  cubic  centimeters 
of  concentrated  H2SO4,  holding  the  tube  in  an  inclined 
position  so  that  the  acid  will  run  down  the  side  to  the 
bottom  without  mixing  with  the  solution.    Allow  the 
tube  to  stand  quietly  for  several  minutes.     If  a  nitrate 
is  present,  a  brown  or  black  ring  will  develop  at  the 
point  where  the  acid  and  the  solution  are  in  contact. 

Explain  this  test  fully  and  write  all  equations.  Obtain 
several  salts  from  the  instructor  and  test  them  to  find 
which  are  nitrates. 

Aqua  Regia. 

155.  Mix  about  i  cc.  concentrated  HNOs  with  3  cc. 
concentrated  HC1  in  a  test  tube  and  warm  the  mixture 
gently.     Is  a  gas  evolved?     If  so,  has  it  any  odor? 
What  is   the   gas?    Explain   the   reaction.     Would   a 


72  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

mixture  of  H2S04  and  HC1  react  in  the  same  way? 
Why? 

156.  Introduce  small  pieces  of  gold  foil  into  each  of 
two  clean  dry  test  tubes.     To  one  add  a  few  drops  of 
concentrated  HNOs  and  to  the  other  a  little  concen- 
trated HC1.     Do  you  notice  any  reaction? 

Heat  the  contents  of  each  tube  to  the  boiling  point. 
Does  this  cause  the  gold  to  dissolve?  While  still  hot, 
pour  the  contents  of  one  tube  into  the  other  and  note  the 
result.  Mention  another  metal  which  dissolves  in  aqua 
regia  but  is  not  soluble  in  either  HC1  or  HN03.  (Label 
the  solution  of  gold  in  aqua  regia  and  save  for  a  future 
experiment.) 

Nitrogen  Iodide  (NTs) . 

157.  In  a  small  beaker  treat  10  cc.  of  tincture  of 
iodine  with  15  cc.  of  strong  NH4OH.     Filter  and  wash 
the  brown  precipitate  on  the  filter.     Tear  the  paper  into 
four  pieces  and  spread  them  on  the  desk  to  dry.     When 
perfectly  dry,  touch  them  one  at  a  time  with  a  glass  rod. 
(CAUTION!) 

Summary.  How  many  oxides  of  nitrogen  are  there? 
Which  is  the  most  stable  oxide  ?  Why  ?  How  many 
acids  of  nitrogen  are  possible?  How  many  of  them  are 
common  ?  In  what  two  ways  does  nitric  acid  act  ? 

Problems,  (a)  What  weight  of  copper  will  be  necessary  to 
produce,  when  treated  with  an  excess  of  HNO3,  1200  liters  of  NO 
at  10°  and  760  mm.  pressure  ? 

(b)  How  many  pounds  of  65%  HNO8  can  be  obtained  from  i 
ton  of  caliche  containing  85%  NaNO3  ? 

(c)  What  weight  of  phosphorus  would  be  necessary  to  burn  all 
the  oxygen  in  i  cubic  meter  of  air  ?    What  volume  of  nitrogen 
at  22°  and  735  mm.  pressure  would  be  left  ? 


CHAPTER  VIII. 
OXIDATION   AND   REDUCTION. 

158.  Burn  a  piece  of  wood.    What  is  the  chief  in- 
gredient in  the  wood  ?    What  becomes  of  it  in  burning  ? 
Is   this   a   case   of   oxidation   or   of   reduction?     How 
does  the  burning  of  wood  compare  with  the  rusting  of 
iron? 

159.  Heat  the  solution  of  gold  chloride  (AuCls)  ob- 
tained in  Exp.  156  to  boiling  for  a  moment.     Cool  the 
solution  under  the  faucet.     Then  add  a  few  cubic  centi- 
meters of  stannous  chloride  (SnCl2)  solution.     Allow  to 
stand  for  a  few  moments  and  notice  the  change.     What 
sort  of  a  change  is  this?     How  does  this  experiment 
compare  with  Exp.  14?    Why  was  the  solution  boiled 
before  adding  the  SnC^  ? 

Does  oxidation  always  mean  the  adding  of  oxygen? 
Does  reduction  always  mean  taking  away  oxygen  ? 

1 60.  To  a  few  cubic  centimeters  of  a  fresh  solution 
of  ferrous  sulphate  (FeSO4)  add  a  few  drops  of  KCNS 
solution.     Does  this  produce  any  change  in  the  appear- 
ance of  the  iron  solution  ? 

To  a  second  portion  of  the  FeS04  solution  add  a  few 
drops  of  concentrated  H2S04  and  a  drop  of  concentrated 
HNO3.  Heat  to  boiling.  Cool  by  holding  under  the 
faucet.  When  cool,  test  with  a  few  drops  of  KCNS 
solution.  Is  there  a  change  this  time?  Explain  why 
these  two  tests  differ.  In  the  second  test,  has  the  iron 
been  oxidized  or  reduced  ?  What  was  the  agent  which 
brought  about  this  oxidation  or  reduction  ? 

73 


74 


EXPERIMENTS  IN   GENERAL  CHEMISTRY 


161.  To  a  few  cubic  centimeters  of  ferric  chloride 
(FeCls)  solution  add  a  drop  of  KCNS.  Compare  with 
the  previous  experiment. 

To  a  second  portion  of  FeCls  solution  add  a  few  cubic 
centimeters  of  SnC^  solution  and  then  test  with  KCNS. 
What  difference  do  you  observe  in  these  two  tests? 
Explain  fully. 

Oxidizincj 


d 

.  R^eduoincj 


FIG.  22.  FIG.  23. 

162.  Oxidizing  and  Reducing  Flame.  Make  a  borax 
bead  in  a  loop  at  the  end  of  a  piece  of  platinum  wire  as 
shown  in  Fig.  22.  To  make  the  bead,  heat  the  wire  to 
redness  and  plunge  into  some  borax.  Again  heat  until 
the  borax  which  clings  to  the  wire  has  melted  and  no 
longer  effervesces.  If  the  loop  is  not  completely  filled 
with  borax,  add  more  in  the  same  way.  Heat  the  bead 
until  it  is  perfectly  clear. 


OXIDATION   AND   REDUCTION 


Pick  up  a  minute  particle  of  Mn02  with  the  hot  bead 
and  then  heat  in  the  hottest  part  of  the  flame  (shown  at 
a  in  Fig.  23)  until  the  bead  is  completely  fused.  Then 
raise  the  bead  in  the  flame  until  it  is  in  the  oxidizing 
flame  as  shown  in  the  drawing.  After  a  few  moments, 
withdraw  the  bead  from  the  flame,  allow  to  cool  and  then 
examine  carefully.  Is  the  bead  colored  ? 


d 


FIG.  24. 


Again  melt  the  bead  in  the  hot  part  of  the  flame  and 
hold  for  several  minutes  in  the  reducing  flame  as  shown 
in  Fig.  23.  Then  lower  it  into  the  green  cone  to  cool. 
Quickly  withdraw  from  the  flame  and  examine.  Has 
the  color  changed  ?  What  is  the  explanation  ? 

Heat  the  bead  in  the  oxidizing  flame  again.  Does  it 
change  to  the  original  color? 

In  a  similar  manner  try  the  action  of  the  oxidizing  and 
the  reducing  flames  on  borax  beads  containing  traces  of 
copper,  cobalt,  iron  and  chromium  compounds  respec- 
tively. Tabulate  results. 

163.  The  Blowpipe.  For  use  with  the  blowpipe,  a 
small  luminous  flame  is  preferable,  as  shown  at  a  in  Fig. 
24.  To  produce  the  reducing  flame,  shown  at  b,  hold 


76  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

the  tip  of  the  blowpipe  just  outside  the  upper  part  of  the 
flame  and  blow  gently  and  evenly  into  it. 

The  oxidizing  flame  is  produced  by  blowing  a  strong 
blast  into  the  flame  as  shown  at  c,  holding  the  tip  of  the 
blowpipe  in  the  lower  part  of  the  flame. 

Make  a  borax  bead  with  Mn02  and  try  the  action  of 
the  oxidizing  and  the  reducing  blowpipe  flames  upon  it. 

164.  Try  the  blowpipe  flames  on  borax  beads  con- 
taining respectively  particles  of  iron  and  copper  com- 
pounds. 

Summary.  Make  a  list  of  all  the  oxidizing  agents  you 
have  studied.  Make  a  similar  list  of  all  the  reducing 
agents  you  have  studied.  Can  you  name  any  other 
oxidizing  agents  or  reducing  agents?  Define  oxidation 
and  reduction. 

What  acid  have  you  studied  which  is  neither  an  oxi- 
dizing nor  a  reducing  agent? 

In  what  does  "combustion"  differ  from  "oxidation"? 
Can  oxidation  take  place  in  solution  ?  Can  combustion  ? 


CHAPTER  IX. 
SULPHUR  (S;  32). 

165.  Place  about  5  gms.  of  roll  sulphur  in  a  test  tube 
and  heat  very  gently  and  gradually.     Notice  the  thin, 
straw-colored  liquid  which  is  formed  when  the  sulphur 
first  melts.     Now  increase  the  heat  gradually  and  note 
all  changes  in  appearance. 

1 66.  To  a  solution  of  calcium  polysulphide  (CaSs)  add 
HC1.     To  a  solution  of  sodium  thiosulphate  (Na2S2O3) 
add  dilute  H2SO4.     What  precipitates  are  formed  in  these 
two  experiments?     Does  it  make  any  difference  what 
acid  is  used  ? 

167.  Put  about  5  gms.  of  sulphur  in  a  test  tube  and 
apply  heat  gently  just  to  melt  the  sulphur.     Pour  the 
thin,  straw-colored  molten  sulphur  into  a  beaker  of  water. 
Examine  the  product. 

In  a  second  test  tube  melt  about  5  gms.  of  sulphur  and 
heat  to  the  boiling  point;  then  pour  into  a  beaker  of 
water.  Examine  the  product  formed.  Does  it  differ 
from  that  formed  in  the  first  part  of  the  experiment? 
Reserve  the  product  for  several  days  and  examine  for 
changes. 

1 68.  Melt  some  sulphur  in  a  small  Hessian  crucible, 
keeping  the  temperature  as  low  as  possible.     Continue 
to  add  sulphur  until  the  crucible  is  full  of  the  liquid. 

Allow  to  cool  until  the  sulphur  begins  to  solidify  on 
the  sides  of  the  crucible ;  then  pour  out  the  molten  sul- 
phur. Carefully  examine  the  crystals  which  line  the 
crucible. 

77 


78  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

169.  Place  about  10  gms.  of  flowers  of  sulphur  in  a 
small  flask  and  add  15  or  20  cc.  of  carbon  disulphide 
(CS2) .     (CAUTION  !  Never  work  with  CS^  when  a  flame  is 
near.)     Shake  for  a  minute  or  two  and  then  filter,  allow- 
ing the  filtrate  to  run  into  a  small  beaker.     Allow  the 
beaker  containing  the  CS2  to  stand  quietly.     When  the 
CS2  is  all  evaporated,  what  is  left  in  the  beaker  ?     Com- 
pare this  product  with  that  obtained  from  Exp.  168  in 
which  sulphur  was  melted  in  a  Hessian  crucible. 

170.  Sulphur    Monochloride     (S2C12).      Arrange    an 
apparatus  as  shown  in  Fig.    25.     The  distilling  flask 
should   contain  about  30  gms.   of  sulphur.     Generate 
chlorine  by  means  of  MnO2,  NaCl  and  H2SO4  and  pass 
the  gas  into  the  distilling  flask  as  shown.     The  second 
distilling  flask,  which  acts  as  a  condenser  and  receiver, 
is  cooled  by  means  of  a  stream  of  water. 

When  the  stream  of  chlorine  becomes  regular,  heat  the 
distilling  flask  to  melt  the  sulphur  and  to  distill  the  S2C12 
formed.  The  thermometer  shows  the  temperature  of  the 
vapor  of  the  latter.  Do  not  let  the  temperature  rise 
above  160°  or  170°. 

When  sufficient  S2C12  has  been  obtained,  stop  the 
operation  and  thoroughly  clean  and  dry  the  distilling 
flask.  Then  introduce  the  S2C12  and  stopper  the  flask 
with  a  cork  carrying  only  a  thermometer.  Connect  the 
apparatus  as  shown  and  redistill  the  liquid,  being  careful 
to  note  the  temperature  at  which  it  distills. 

171.  Describe  fully  the  properties  of  sulphur  mono- 
chloride.     What  sort  of  an  odor  has  it?     Drop  a  little 
of  the  liquid  into  water  and  note  the  effect.     To  a  few 
drops  of  linseed  oil  in  a  test  tube  add  a  few  drops  of 
S2C12.    What  happens? 


SULPHUR 


79 


Summary.  How  many  varieties  of  sulphur  are  there  ? 
What  name  do  we  apply  to  various  forms  of  the  same 
element?  What  other  element  have  we  already  studied 
which  has  more  than  one  form  ?  What  is  the  difference 
between  "roll  sulphur"  and  " flowers  of  sulphur"? 


FIG.  25. 

Hydrogen  Sulphide  (H2S) . 

172.  Add  a  few  drops  of  HC1  to  a  small  piece  of  iron 
sulphide   (FeS)   in   an   evaporating  dish  or  test  tube. 
Does  the  gas  which  is  formed  have  an  odor? 

173.  Arrange  an  apparatus  for  generating  H2S  similar 
to  that  used  for  hydrogen  (Fig.  3,  page   7).      In  the 
flask  place  a  few  lumps  of  FeS  and  through  the  funnel 


80  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

gradually  add  dilute  HC1  until  a  regular  evolution  of  the 
gas  results. 

Collect  a  bottle  of  the  gas  by  displacement  of  air  and 
apply  a  match.  Does  the  gas  burn  or  support  combus- 
tion ?  What  are  the  products  formed  ? 

174.  Pass  a  stream  of  H2S  for  several  minutes  through 
50  cc.  of  water  in  a  loo-cc.  flask,  occasionally  shaking  the 
flask  to  hasten  absorption.     Try  the  action  of  this  solu- 
tion  on   a   few  cubic   centimeters  of  copper   sulphate 
(CuSO4)    solution.     In   a   second   test   tube  pass  H2S 
through  a  few  cubic  centimeters  of  CuSO4  solution.     Is 
the  action  the  same  with  the  gas  as  with  the  solution  of 
the  gas  in  water? 

Try  the  action  of  the  H2S  solution  on  litmus  paper. 
What  do  you  consider  to  be  the  nature  of  the  compound 
H2S? 

175.  In  separate  test  tubes  try  the  action  of  H2S  on 
solutions  of  any  compounds  of  silver,   mercury,   lead, 
cadmium,  bismuth,  copper,  arsenic,  antimony  and  tin. 
Express  the  results  in  the  form  of  a  table. 

176.  Pass  H2S  through  a  solution  of  NiSO4  for  a  few 
moments;  then  add  NH4OH.     Does  the  latter  cause  any 
change  to  take  place?     Why? 

Make  a  solution  of  ammonium  sulphide  ((NH4)2S)  by 
passing  H2S  gas  through  50  cc.  of  dilute  NH4OH.  Try 
the  action  of  this  solution  on  solutions  of  any  com- 
pounds of  iron,  cobalt,  nickel,  manganese  and  zinc. 
Express  all  results  in  the  form  of  a  table  as  in  the  pre- 
vious experiment. 

177.  Pass  H2S  through  a  few  cubic  centimeters  of 
concentrated  HN03  in  a  small  flask.     What  happens? 
Can  you  write  the  equation  ? 


SULPHUR  8 1 

Pass  H2S  through  a  few  cubic  centimeters  of  concen- 
trated H2S04  in  a  small  flask.    Write  all  equations. 
What  sort  of  an  action  has  H2S  in  these  two  cases  ? 

178.  Try  the  action  of  H2S  gas  on  a  few  cubic  centi- 
meters of  potassium  dichromate  (K2Cr2O7)  solution  con- 
taining a  few  drops  of  concentrated  HC1.     Repeat,  using 
potassium  permanganate  (KMn04)  solution  instead  of 
K2Cr207. 

Summary.  In  what  two  ways  does  H2S  act?  What 
two  groups  of  metals  cannot  be  precipitated  by  means  of 
this  compound  nor  by  (NH4)2S  solution  ?  What  elements 
are  precipitated  as  sulphides  by  (NH4)2S  but  not  by 
H2S  ?  Which  sulphide  is  white  ? 

Sulphur  Dioxide  (SO2),  Sulphurous  Acid  (H2S03). 

179.  Burn  a  small  piece  of  sulphur  and  note  the  odor 
of  the  gas  formed.     In  a  small  porcelain  crucible,  strongly 
heat  a  small  piece  of  iron  pyrites  (FeS2)  and  note  the 
odor. 

1 80.  By  means  of  a  deflagrating  spoon,  burn  a  small 
piece  of  sulphur  in  a  5oo-cc.  flask  containing  about  20  cc. 
of  water.     Cork  the  flask  and  shake.     Test  the  action 
of  the  water  on  litmus  paper.     What  has  been  formed 
and  is  now  in  solution  in  the  water? 

181.  In  a  test  tube  heat  a  small  piece  of  metallic 
copper  with  a  few  cubic  centimeters  of  concentrated 
H2S04.    What  gas  is  evolved?     Repeat,  using  a  small 
piece  of  charcoal  (C)  instead  of  the  copper. 

What  sort  of  an  action  have  Cu  and  C  on  hot 
H2S04? 

182.  For  generating  SO2  arrange  an  apparatus  similar 
to  that  used  for  the  generation  of  chlorine  (Fig.  n, 


82  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

page  36).  Place  about  25  gms.  of  copper  turnings  in 
the  flask  and  add  about  30  cc.  of  concentrated  H2SO4. 

Gently  warm  the  flask  and  when  the  gas  (802)  begins 
to  be  evolved  in  a  regular  stream,  collect  one  or  two 
bottles  by  displacement  of  air  and  cover  with  glass 
plates. 

Now  prepare  an  aqueous  solution  of  the  gas  by  pass- 
ing the  latter  through  about  50  cc.  of  water  in  a  zoo-cc. 
flask,  until  the  water  is  saturated. 

183.  Into  one  of  the  bottles  of  S02  prepared  above 
thrust  a  burning  splinter.     What  happens?     Is  the  gas 
a  supporter  or  a  non-supporter  of  combustion?     Does 
it  burn?     Introduce  a  piece  of  wet  blue  litmus  paper 
into  one  of  the  bottles  of  the  gas.    Dip  a  piece  of  blue 
litmv~  paper  into  the  aqueous  solution  of  SO2.     What 
sort  of  a  reaction  do  you  get  with  the  litmus  ?    What  is 
formed  by  the  combination  of  S02  and  water? 

184.  In  separate  test  tubes  try  the  action  of  the 
sulphurous  acid  (H2S03)  made  in  Exp.  182  on  solutions 
of  the  following  substances:    potassium  permanganate 
(KMnO4),  copper  sulphate  (CuS04),  potassium  dichro- 
mate  (K2Cr207),  litmus,  indigo,  and  cochineal. 

What  sort  of  substances  are  bleached  by  sulphurous 
acid?  What  element  have  we  previously  studied  which 
acts  in  the  same  way  ? 

Sulphur  trioxide  (SOs),  Sulphuric  acid  (H2SO4). 

185.  To  prepare  H2SO4  by  the  "lead  chamber"  proc- 
ess, construct  the  apparatus  shown  in  Fig.  26.    Water 
is  boiled  in  one  of  the  small  flasks,  NO  generated  in 
another,  and  S02  in  another.    The  steam,  S02  and  NO 
are  conducted  into  the  large  flask  where  they  come  to- 


SULPHUR  83 

gether  with  the  formation  of  H2SO4.  It  is  not  necessary 
to  continue  the  generation  of  NO  for  any  length  of  time 
inasmuch  as  little  is  lost.  Air  should  be  blown  from 
time  to  time  through  the  rubber  tube  shown  in  the 
illustration. 


FIG.  26. 

Discontinue  the  production  of  steam  for  a  moment 
and  notice  the  formation  of  " chamber  crystals"  on  the 
inside  walls  of  the  flask.  What  is  the  composition  of 
these  crystals  ? 

If  air  is  not  blown  into  the  flask  for  a  few  moments,  the 
contents  will  become  colorless.  Then,  by  blowing  air 
into  the  apparatus,  the  brown  color  of  N02  is  produced. 

After  running  the  experiment  for  some  moments,  dis- 
connect the  apparatus  and  reserve  the  H2S04  which 
has  been  formed  for  use  in  Exp.  186. 

1 86.  Try  the  action  of  dilute  H2SO4  from  the  reagent 
bottle  on  solutions  of  the  following  substances:  barium 
chloride  (BaCl2),  strontium  chloride  (SrCl2),  calcium 
chloride  (CaCl2),  sodium  chloride  (NaCl)  and  lead  acetate 


84  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

(Pb(C2H302)2).     Which  metals  form  insoluble  sulphates? 
Repeat  the  tests,  using  the  H2S04  prepared  in  Exp.  185. 

187.  To  a  few  cubic  centimeters  of  a  saturated  solution 
of  sugar,  in  a  large  beaker,  add  a  few  cubic  centimeters 
of  concentrated  H2SO4.     What  happens  ? 

To  a  small  splinter  of  wood  in  a  beaker  or  on  a  watclj. 
glass  add  a  few  cubic  centimeters  of  concentrated  H2SO4. 
Is  the  action  similar  to  that  with  sugar  ? 

188.  By   means   of   a   graduated   cylinder   carefully 
measure  36  cc.  of  water  and  pour  it  into  a  large  dry 
beaker.     Dry  the  graduate  as  well  as  possible  and  care- 
fully measure  out  53  cc.  of  concentrated  H2SO4.     Slowly 
pour  the  acid  into  the  water  in  a  fine  stream.     Does  the 
beaker  become  hot  ?    Why  ?    Allow  the  mixture  to  cool; 
then  carefully  measure  its  volume.     Has  there  been  a 
contraction  or  an  expansion?    What  has   caused  it? 
(Pour  the  diluted  acid  into  the  large  stock  bottle  labeled 
" Dilute  H2SO4.") 

189.  Place  one  or  two  drops  of  concentrated  H2SC>4  in 
a  clean  dry  evaporating  dish  and  heat  strongly.     What 
happens?     (Avoid  breathing  the  fumes.) 

190.  Make  a  mixture  of  powdered  BaS04  and  Na2CO3 ; 
place  the  mixture  on  a  piece  of  charcoal  and  heat  strongly 
with  the  blowpipe  flame.     When  cold,  place  the  fused 
mass  on  a  clean  silver  coin  and  add  a  drop  or  two  of  water. 
Does  the  coin  become  stained?    Why? 

Repeat  the  experiment,  using  some  other  sulphate. 

191.  In  separate  test  tubes  try  the  action  of  BaCl2 
solution  on  solutions  of  sodium  sulphate  (Na2SO4),  cop- 
per sulphate  (CuS04),  and  magnesium  sulphate  (MgSO4). 
Does  BaCl2  react  the  same  with  soluble  sulphates  as 
withH2S04? 


SULPHUR  85 

Summary.  How  many  series  of  salts  has  H2S04  ? 
(See  Exp.  1 16.)  Write  the  structural  formula  for  H2SO4. 
What  is  the  test  for  sulphuric  acid  and  soluble  sulphates  ? 
For  insoluble  sulphates?  For  a  sulphide?  For  free 
sulphur  ? 

What  is  pyro-sulphuric  acid?  Why  is  it  a  strong 
chemical  ? 

Thio-sulphuric  Acid  (H2S2O3). 

192.  To  about  50  cc.  of  sodium  sulphite  (Na2SO3) 
solution  in  a  beaker  add  about  5  gms.  of  flowers  of 
sulphur  and  boil  for  a  few  minutes.     Filter;  to  a  portion 
of  the  clear  filtrate  add  a  few  drops  of  concentrated 
H2S04  and  allow  to  stand  for  a  time.     To  a  second  por- 
tion add  a  little  tincture  of  iodine. 

In  separate  test  tubes  try  the  action  of  a  few  drops  of 
concentrated  H2S04  and  a  few  drops  of  iodine  solution  on 
a  solution  of  sodium  sulphite  (Na2SO3).  Do  these  tests 
give  the  same  results  as  those  performed  previously  ?, 

193.  To  a  solution  of  AgNO3  add  Na2S2O3  solution, 
a  drop  at  a  time,  until  the  precipitate  at  first  formed 
redissolves. 

194.  Prepare  AgCl  by  adding  an  NaCl  solution  to  a 
few  cubic  centimeters  of  a  solution  of  AgNU3.     Try  the 
action  of  Na2S2O3  on  the  precipitate.     What  commercial 
applications  are  made  of  this  reaction  ? 

195.  In  a  test  tube  try  the  action  of  Na2S2O3  solution 
on  a  few  cubic  centimeters  of  a  solution  of  KMnO4. 

Repeat,  using  iodine  solution  instead  of  KMnO4. 
What  other  oxygen  acid  of  sulphur  is  produced  by  this 
latter  reaction  ? 

Summary.  What  common  use  is  made  of  Na2S203? 
What  are  the  common,  though  incorrect,  names  for  this 


86  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

compound?  Why  is  the  salt  called  a  " thio-sulphate " ? 
If  it  were  possible  to  make  such  a  compound  as  a  "thio- 
nitrate,"  what  would  be  the  formula  for  it? 

Problems,  (a)  What  volume  of  H2S  gas,  at  25°  and  860  mm., 
can  be  obtained  by  the  action  of  an  excess  of  acid  on  3800  gms. 
of  72%  Sb2S3? 

(b)  By  roasting  20  tons  of  90%  ZnS  ore  and  using  the  by-prod- 
uct in  the  manufacture  of  H2SO4,  what  volume  of  45%  H2SO4  can 
theoretically  be  obtained? 

(c)  What  weight  of  precipitated  sulphur  can  be  obtained  by 
treating  600  gms.  of  Na^Os,  dissolved  in  water,  with  excess  of 
H2S04? 

(d)  What  weight  of  gas  (o°  and  760  mm.)  will  be  obtained  by 
heating  together  36  gms.  of  pure  carbon  and  an  excess  of  con- 
centrated H2SO*? 


CHAPTER  X. 
CARBON  (C;  12). 

196.  Place  a  small  piece  of  wood  in  an  iron  crucible 
and  cover  with  a  layer  of  sand.     Heat  with  the  full  force 
of  the  Bunsen  burner  until  gas  no  longer  comes  from  the 
crucible  and  burns.     Allow  to  cool.     Examine  the  con- 
tents of  the  crucible.     What  is  the  composition  of  wood  ? 

197.  Place  about  5  gms.  of  cane  sugar  in  a  porcelain 
evaporating  dish  and  heat  gently  until  there  is  no  further 
change  in  the  appearance  of  the  material  in  the  dish. 
What  are   the   chief  products  formed  when   sugar  is 
heated  ? 

198.  Turn  off  the  air  supply  on  a  Bunsen  burner  so 
that  it  burns  with  a  luminous  flame.     Hold  a  piece  of 
cold  porcelain  (crucible,  evaporating  dish  or  mortar)  in 
this  flame  and  observe  the  black  deposit. 

Compare  the  products  formed  in  these  three  experi- 
ments. Compare  them  with  the  bone  black  on  the  side 
shelf.  Why  do  we  not  make  bone  black  in  the  labora- 
tory? 

199.  Fill  a  hard  glass  tube  with  small  pieces  of  soft 
coal.     Connect  a  delivery  tube  as  shown  in  Fig.  6,  page 
15.    Heat  the  tube  strongly  and  collect  the  gas  evolved 
in  bottles  by  displacement  of  water.     Examine  this  gas. 
Will  it  burn  ?    Has  it  any  odor  ? 

Examine  the  residue  in  the  tube  and  compare  it  with 
the  other  varieties  of  carbon  formed  in  the  preceding 
experiments.  Notice  the  water  over  which  the  gas  was 
collected. 

87 


88  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

200.  Heat  a  bit  of  charcoal  on  platinum  foil.     Like- 
wise heat  a  bit  of  graphite  on  platinum  foil.     Do  both 
substances  burn?     Which  is  the  more  stable  form  of 
carbon?     Name  another  stable  variety  of  carbon. 

201.  Add  about  a  gram  of  powdered  bone  black  to 
15  cc.  of  an  indigo  solution.     Boil  for  a  few  minutes. 
Filter  and  examine  the  filtrate. 

Repeat,  using  litmus  solution  instead  of  indigo. 

202.  Into  a  small  flask  containing  15  or  20  cc.  of  H2S 
solution  introduce  about  2  gms.  of  powdered  bone  black. 
Cork  the  flask  and  shake  for  several  minutes.     After 
standing  for  15  minutes,  filter  the  contents  and  test  a 
portion  of  the  clear  filtrate  for  H2S  by  adding  a  few  drops 
of  lead  acetate  (Pb(C2H3O2)2)  solution. 

What  commercial  use  is  made  of  bone  black? 

APPROXIMATE  ANALYSIS  OF  COAL. 

(Quantitative.) 

203.  Determination    of    Volatile    Matter.     Carefully 
weigh  a  clean  dry  porcelain  crucible  and  cover.     Intro- 
duce about  2  gms.  of  powdered  coal  into  the  crucible  and 
again  weigh  carefully.     The  difference  in  weight  repre- 
sents the  weight  of  the  coal  taken. 

Place  the  crucible,  covered,  on  a  triangle  on  a  ring- 
stand  and  heat  strongly  until  gases  no  longer  come  from 
the  crucible  and  burn.  Allow  to  cool;  then  weigh  care- 
fully. 

Calculate  the  loss  in  weight  of  the  coal.  This  loss  in 
weight  represents  volatile  matter  in  the  coal.  Express 
your  result  as  "percentage  of  volatile  matter." 

Examine  the  residue  in  the  crucible.  What  is  the 
nature  of  it  ? 


CARBON  89 

Determination  of  Ash.  Carefully  weigh  a  clean  dry 
porcelain  crucible  without  the  cover.  Introduce  about 
2  gms.  of  powdered  coal  and  again  weigh  accurately. 
Place  the  crucible  on  a  triangle  and  cover  with  a  crucible 
cover.  Apply  heat  with  the  crucible  covered  until  gases 
no  longer  come  from  the  crucible  and  burn.  Then  re- 
move the  cover  and  place  the  crucible  in  an  inclined 
position  on  the  triangle  so  that  air  may  enter  freely. 
Heat  with  the  full  force  of  the  Bunsen  burner  until  all 
carbon  is  burned  and  the  ash  in  the  crucible  is  white  or 
gray  in  color. 

Allow  the  crucible  to  cool;  then  weigh.  Calculate  the 
percentage  of  ash  in  the  coal. 

Assuming  that  the  determinations  of  ash  and  volatile 
matter  are  accurate,  what  is  the  percentage  of  carbon  in 
the  sample  of  coal  used?  Record  the  number  of  the 
sample  of  coal  analyzed. 

Carbon  Monoxide  (CO). 

204.  Prepare  carbon  monoxide  (GO)  by  heating  to- 
gether 15  gms.  of  oxalic  acid  (HzCzO*)  and  50  cc.  of  con- 
centrated H2S04  in  an  apparatus  as  shown  in  Fig.  27. 
The  wash  bottle  should  contain  a  concentrated  solution 
of  NaOH. 

Collect  several  bottles  of  the  gas  over  water.  Ascer- 
tain if  it  will  burn  or  support  combustion.  Add  a  few 
cubic  centimeters  of  lime  water  or  baryta  water  to  the 
bottle  tested;  cover  and  shake. 

205.  Disconnect  the  apparatus  at  point  a  and  in  place 
of  the  delivery  tube  connect  a  piece  of  hard  glass  tubing 
containing  a  little  black  copper  oxide.    While  continu- 
ing the  generation  of  CO,  strongly  heat  the  copper  oxide 


90  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

in  the  glass  tube.     What  happens  ?    Pass  the  gas  coming 
from  the  apparatus  into  lime  water. 

206.  Place  a  few  crystals  of  oxalic  acid  in  a  test  tube 
and  add  a  little  concentrated  H2SO4.  Heat  until  re- 
action takes  place.  Test  the  gas  evolved  (i)  for  CO,  by 
bringing  the  mouth  of  the  test  tube  to  a  flame,  and  (2) 
for  C02  by  holding  a  glass  rod  with  a  drop  of  lime  water 
at  the  end  in  the  mouth  of  the  tube. 


FIG.  27. 

207.  Arrange  an  apparatus  as  shown  in  Fig.  28,  con- 
sisting of  CO2  generator,  wash  bottle  and  hard  glass  tube. 
Partially  fill  the  hard  glass  tube  with  powdered  zinc. 
Then  generate  CO2  in  the  flask  by  means  of  marble  and 
HC1,  and  strongly  heat  the  hard  glass  tube  containing 
the  zinc.     Test  the  gas  coming  from  the  end  of  the  hard 
glass  tube  to  ascertain  if  it  will  burn. 

208.  Introduce  an  intimate  mixture  of  about  two 
parts  by  weight  of  powdered  CaCOs  and  one  part  of 


CARBON 


powdered  zinc  into  a  hard  glass  test  tube  or  a  piece  of 
hard  glass  tubing  closed  at  one  end.  Arrange  a  delivery 
tube  as  shown  in  Fig.  6,  page  15.  Strongly  ignite  and 
then  collect  one  or  two  bottles  of  the  gas  which  is  evolved. 
Test  the  gas.  What  is  it? 


FIG.  28. 

Carbon  Dioxide  (C02),  Carbonic  Acid  ( 

209.  By  means  of  a  piece  of  glass  tubing,  blow  through 
a  little  lime  water  in  a  beaker  for  several  minutes.     What 
causes  the  change?     Compare  with  Exp.  122,  page  59.  -J 

210.  Place  a  small  piece  of  marble  (CaCO3)  in  a  test 
tube  and  add  HC1.     Test  the  gas  evolved  by  holding  a 
glass  rod  with  a  drop  of  lime  water  at  the  end  in  the 
mouth  of  the  tube. 

211.  Arrange  an  apparatus  consisting  of  a  CO2  gener- 
ator .and  wash  bottle  nearly  filled  with  water  as  shown 
in  Fig,  29.     Place  a  few  lumps  of  marble  in  the  flask, 
cover  with  water  and  add  concentrated  HC1  through 


92 


EXPERIMENTS   IN   GENERAL  CHEMISTRY 


the  thistle  tube  to  produce  a  brisk  evolution  of  the 
gas.  Collect  several  bottles  of  C02  over  water.  Cover 
with  glass  plates  and  reserve  for  use  in  Exps.  215 
and  216. 

212.  Change  the  delivery  tube  at  a  as  shown  by  the 
dotted  line  and  pass  C02  through  lime  water  until 
there  is  no  further  change.  Be  careful  to  notice  all 
changes. 


FIG.  29. 

213.  Pass  CO2  through  100  cc.  of  water  for  several 
minutes.     Taste   the   water.     Try   its   action   towards 
litmus  paper.     What  is  in  solution  in  the  water? 

214.  Pass  C02  through  about  25  cc.  of  dilute  NaOH 
solution  until  the  latter  is  saturated.     What  is  formed  ? 
Add  dilute  HC1  to  the  solution.     Does  this  prove  the 
nature  of  the  compound  in  solution?     Compare  with 
Exp.  210. 

215.  Into  one  of  the  bottles  of  C02  collected  in^Exp. 
211  introduce  a  piece  of  wet  litmus  paper.     To  another 
add  a  few  drops  of  baryta  or  lime  water.     Plunge  a 


CARBON  93 

burning  splinter  into  another  bottle  of  the  gas.  Into  a 
fourth  bottle  introduce,  by  means  of  forceps,  a  piece  of 
burning  magnesium  ribbon  which  has  been  previously 
cleaned  with  sand  paper. 

216.  Ignite  a  few  drops  of  benzene  on  a  watch  glass. 
Pour  CO2  from  a  bottle  onto  the  burning  benzene.  What 
happens  ?  What  commercial  application  is  made  of  this 
reaction  ? 


DETERMINATION  OF  THE  WEIGHT  OF  A  LITER  OF 
CARBON  DIOXIDE. 

(Quantitative.) 

217.  This  experiment  is  carried  out  in  exactly  the 
same  manner  as  the  determination  of  the  weight  of  a 
liter  of  chlorine  (Exp.  63,  page  38).     The  CO2  may  be 
generated  by  the  action  of  HC1  on  CaCO3  and  must  be 
passed  through  a  wash  bottle  containing  water,  and  a 
drying  bottle  containing  concentrated  H2S04.     Why  are 
the  wash  bottle  and  the  drying  bottle  necessary  ? 

Carbon  Bisulphide  (CS2). 

218.  Place  a  few  drops  of  CS2  on  the  hand  and  blow 
across  it.     Note  the  color  and  odor  of  the  liquid.     Will 
it  mix  with  water  ? 

219.  Pour  a  few  drops  of  CS2  into  a  porcelain  dish. 
Heat  a  glass  rod  to  redness  and  hold  over  the  dish  just 
above  the  liquid.     Note  the  products  formed  when  CS2 
burns.     How  many  products  can  you  identify?    Write 
several  equations  to  represent  the  burning  of  €82  with 
varying  amounts  of  oxygen. 


94          EXPERIMENTS  IN  GENERAL  CHEMISTRY 

DETERMINATION  OF  SPECIFIC  GRAVITY  OF  CARBON 
BISULPHIDE. 

220.  For  this  experiment  a  small  flask  with  a  tightly 
fitting  cork  or  rubber  stopper  may  be  employed.  With 
a  triangular  file  make  a  scratch  about  half  way  up  the 
neck  of  the  flask. 

Carefully  clean  the  flask  and  stopper  and  when  both 
are  thoroughly  dry,  insert  the  stopper  in  the  flask  and 
weigh.  Let  this  weight  of  the  empty  flask  be  repre- 
sented by  E.  Now  fill  the  flask  to  the  mark  on  the 
neck  with  distilled  water,  cork  and  weigh.  Let  this 
weight  be  represented  by  W.  Empty  the  flask  and 
rinse  about  four  times  with  small  amounts  of  alcohol 
and  thoroughly  dry  by  means  of  a  blast  of  air.  Fill  the 
flask  to  the  mark  on  the  neck  with  the  sample  of  CS2 
to  be  determined,  cork  and  weigh.  Let  this  weight  be 
represented  by  L. 

W  -  E  equals  the  weight  of  the  water  and  L  —  E 
then  equals  the  weight  of  an  equivalent  volume  of  the 
CS2  being  determined.  The  specific  gravity  of  a  sub- 
stance is  the  relation  of  the  weight  of  a  given  volume  of 
the  substance  as  compared  with  the  weight  of  an  equal 
volume  of  water,  and  inasmuch  as  the  unit  volume  of 
water,  i  cc.,  weighs  i  gram,  the  specific  gravity  of  any 
substance  is  numerically  the  weight  of  a  cubic  centimeter 
of  that  substance  expressed  in  grams.  From  this 

L-E 


Specific  Gravity  = 


W  -E 


Problems,     (a)  How  many  grams  of  95%  CaCO3  will  be  re- 
quired to  produce  1800  liters  of  CO2  at  standard  conditions? 
(b)  What  volume  of  oxygen  will  be  required  to  completely 


CARBON  95 

burn  780  liters  of  CO?    What  volume  of  C02  will  be  formed  (o° 
and  760  mm )  ? 

(c)  How  many  cubic  centimeters  of  CS2  can  be  theoretically 
produced  from  1 20  gms.  of  sulphur  ? 

(d)  If  284  cc.  of  CS2  are  volatilized  at  35°  and  722  mm.  pres- 
sure, what  volume  of  gas  will  be  formed? 

Cyanogen  and  the  Cyanides. 

221.   Fit  a  test  tube  with  a  one-hole  stopper  carrying 
a  short  piece  of  glass  tubing  drawn  to  a  point.     Into 


FIG.  30. 

the  tube  introduce  a  few  grams  of  mercuric  cyanide 
(Hg(CN)2)  and  heat  gently,  letting  the  outlet  tube 
touch  a  second  flame.  (Fig.  30.)  Do  not  breathe  the 
gas.  Why?  (Hood.) 

To  what  is  the  gas  cyanogen,  C2N2,  equivalent 
chemically  ? 

222.  Add  a  solution  of  potassium  cyanide  (KCN),  a 
drop  at  a  time,  until  in  excess,  to  a  solution  of  silver 
nitrate  (AgNO3). 


96  EXPERIMENTS  IN  GENERAL  CHEMISTRY 

223.  Add  potassium  sulphocyanate  solution  (KCNS) 
to  a  little  dilute  FeCls  solution. 

224.  Mix  a  few  drops  each  of  KCN  solution  and 
(NH4)2SX  on  a  watch  glass  and  cautiously  evaporate  to  a 
small  volume.     Test  the  resulting  solution  by  adding  a 
drop  of  FeCls  solution.     Compare  with  previous  experi- 
ment. 

225.  In  a  porcelain  crucible  heat  a  small  piece  of 
KCN  with  about  twice  as  much  lead  oxide   (PbO). 
(HOOD.)     When  entirely  melted,  allow  to  cool;  then  ex- 
tract with  water  and  examine  the  solid  product  formed. 
What  compound  is  in  solution  ? 

226.  Mix  about  15  cc.  each  of  NaOH  and  FeS04 
solution  in  a  small  flask.     Heat  to  boiling,  add  about  15 
cc.  of  KCN  solution  and  boil  for  several  minutes.     Dilute 
with  an  equal  volume  of  water,  filter  and  add  HC1  until 
neutral. 

In  separate  test  tubes  try  the  action  of  small  portions  of 
the  solution  thus  formed  on  solutions  of  CuSO4  and  FeCl3. 

227.  To  the  remainder  of  the  solution  formed  in  the 
first  part  of  Exp.  226,  add  about  20  cc.  of  strong  bromine 
water  and  heat  to  boiling.     Cool.     Test  a  small  portion 
of  FeCls  solution  with  the  solution  thus  formed.     Also 
try  the  action  of  the  solution  on  CuSO4  solution. 

228.  Test  the  action  of  a  solution  of  potassium  ferro- 
cyanide  (K4Fe(CN)6)  on  solutions  of  (i)  FeCla  and  (2) 
FeSO4.     Likewise  try  the  action  of  a  solution  of  potas- 
sium ferricyanide  (K3Fe(CN)e)  on  solutions  of  FeCls  and 
FeS04. 

Express  the  results  of  these  four  tests  in  the  form  of 
a  table,  stating  whether  or  not  a  precipitate  is  formed 
and  also  the  color  produced. 


CARBON  97 

Organic  Chemistry. 

229.  Methane  (CH4).     Into  a  hard  glass  tube,  sealed 
at  one  end,  introduce  an  intimate  mixture  of  equal  parts 
of  fused  sodium  acetate  and  dry  soda  lime.    Arrange 
a  delivery  tube  as  shown  in  Fig.  6,  page  15. 

Strongly  ignite  the  tube  and  collect  the  gas  evolved 
by  displacement  of  water.  Test  the  gas  with  a  burning 
splinter.  Has  the  gas  any  odor  ? 

230.  Acetylene  (C2H2).    Treat  a  few  pieces  of  calcium 
carbide  (CaC2)  with  water  in  a  large  beaker  and  ignite 
the  gas  which  is  produced.     Why  does  it  burn  with  a 
smoky  flame  ?    How  could  you  arrange  an  apparatus  to 
burn  the  gas  with  a  brilliant  smokeless  flame  ? 

231.  Fermentation.     In    a    large    flask    treat    about 
25  cc.  of  molasses  with  200  cc.  of  water.     Add  a  little 
yeast  and  stopper  the  flask  loosely  with  a  wad  of  cotton. 
Allow  it  to  stand  several  days,  preferably  in  a  warm  place. 

Warm  the  flask  and  contents  and  test  for  C02  by  in- 
troducing a  glass  rod  with  a  drop  of  lime  water  or  baryta 
water  on  the  end. 

Filter  the  contents  of  the  flask  and  distill  the  clear 
filtrate  (Fig.  8,  page  23),  collecting  only  the  portion 
which  distills  below  90°.  Examine  this  distillate. 
Notice  the  odor.  What  is  the  compound  ? 

Summary.  What  is  a  good  definition  of  "  organic  " 
chemistry  ? 

What  is  a  hydrocarbon?  What  is  the  name  of  the 
most  simple  series  of  hydrocarbons?  Give  general  for- 
mulas of  the  more  common  series. 

What  is  the  characteristic  group  of  an  alcohol?  Of 
an  organic  acid?  Of  an  aldehyde?  What  general  name 
is  applied  to  organic  salts  ? 


CHAPTER  XI. 

SILICON  AND  BORON. 

SILICON   (Si;  28). 

232.  To  a  few  cubic  centimeters  of  sodium  silicate 
(Na2Si03)  solution  in  a  test  tube  add  a  little  concentrated 
HC1.     What  happens?    Test  the  solubility  of  the  prod- 
uct formed  in  both  acids  and  alkalies. 

Dilute  a  few  cubic  centimeters  of  Na2Si03  solution 
with  three  times  its  volume  of  water  and  add  a  little  con- 
centrated HC1  to  the  mixture.  Is  there  a  precipitate 
formed?  Allow  the  test  tube  containing  the  mixture 
to  stand  quietly  for  a  time.  Does  it  change  in  appear- 
ance? 

233.  Try  the  action  of  a  solution  of  NH4C1  on  a  solu- 
tion of  Na2Si03.     Compare  the  product  with  that  ob- 
tained from  the  first  part  of  Exp.  232.     Explain  fully 
and  write  all  equations. 

234.  In  an  evaporating  dish  add  concentrated  HC1  to 
about  10  cc.  of  Na2SiOa  solution  and  evaporate  to  dry- 
ness.     Treat  the  residue  with  a  little  water.     Filter. 
Examine  the  residue  left  on  the  filter.     Is  it  gelatinous  ? 
Why? 

Dry  the  residue  by  pressing  it  between  pieces  of  filter 
paper  and  save  for  use  in  Exps.  235  and  236. 

235.  In  a  loop  at  the  end  of  a  piece  of  platinum  wire 
make  a  bead  of  "microcosmic  salt"  (NaNH4HPO4)  in 
the  same  way  in  which  borax  beads  are  made  (see  Exp. 
162). 

98 


SILICON  99 

Introduce  a  little  of  the  powdered  silica  (Si02)  from 
Exp.  234  into  the  bead  and  fuse  in  the  Bunsen  flame. 
Observe  the  bead  both  while  in  a  molten  condition  and 
after  cooling.  Is  silica  soluble  in  the  microcosmic  salt 
bead? 

Repeat  this  bead  test  with  a  sample  of  powdered 
quartz  or  sand.  Try  the  test  with  some  insoluble  sili- 
cate in  powdered  condition.  For  what  is  the  bead  a 
good  test  ? 

236.  Into  a  test  tube  introduce  a  mixture  of  a  little 
powdered  quartz  or  sand  and  a  little  fluorspar  (CaF2). 
Moisten  the  mixture  with  concentrated  H2SO4. 

Heat  the  mixture  gently,  at  the  same  time  holding  in 
the  mouth  of  the  test  tube  a  glass  rod  with  a  drop  of 
water  on  the  end. 

Repeat  the  test,  using  Si02  from  Exp.  234.  Repeat 
with  some  insoluble  silicate. 

237.  Make  an  intimate  mixture  of  one  part  SiO2  and 
five  parts  dry  Na2CO3.     Place  the  mixture  in  an  iron 
crucible,  cover   and  heat  strongly.     Occasionally  raise 
the  cover  and  note  the  appearance  of  the  contents.     The 
heating  should  be  continued  until  the  contents  of  the 
crucible  are  in  a  state  of  quiet  fusion. 

Cool,  extract  the  melt  with  water  and  filter.  What 
compound  does  the  filtrate  contain?  Evaporate  a  por- 
tion of  the  filtrate  to  dryness  and  examine  the  residue. 
Test  another  portion  with  concentrated  HC1,  and  a  third 
portion  with  NH4C1  solution.  Compare  with  Exps.  232 
and  233. 

238.  In  separate  test  tubes  try  the  action  of  Na2Si03 
solution   on   solutions   of    CuSO4,    BaCl2,    ZnS04   and 
Pb(C2H,02)2. 


100 


EXPERIMENTS  IN  GENERAL  CHEMISTRY 


Observe  the  nature  of  the  products.    Are  these  in- 
soluble silicates  varieties  of  glass  ?    Why  ? 

239.  Break  a  clean  dry  test  tube,  place  the  pieces  in  a 
clean  dry  porcelain  mortar  and  grind  to  a  fine  powder. 
(CAUTION.    Protect  the  eyes.)     Moisten  the  powder  with 
a  few  drops  of  water  and  then  add  a  drop  of  phenol- 
phthalein.     Explain  the  phenomenon  observed.     What 
are  the  ingredients  of  Bohemian  glass?    What  can  you 
say  of  the  solubility  of  glass  ? 

240.  Fluosilicic  Acid    (H2SiF6).     Arrange  an   appa- 
ratus as  shown  in  Fig.  31.    The  delivery  tube  should  dip 

below  the  surface  of  a 
layer  of  mercury  at  the 
bottom  of  the  cylinder 
of  water.  Use  a  250-0:. 
flask.  It  is  very  neces- 
sary that  the  flask  and 
delivery  tube  be  thor- 
oughly dry  inside,  and 
for  this  reason  the  water 
should  not  be  intro- 
duced into  the  cylinder 
X]  until  the  experiment  is 
started. 

Into  the  flask  intro- 
duce an  intimate  mix- 
ture of  15  gms.  of  fine 
sand  or  powdered  quartz, 
and  10  gms.  of  pow- 


FIG.  31. 


dered  fluorspar  (CaF2).  Through  the  thistle  tube  add 
enough  concentrated  H2S04  to  form  a  thick  paste.  Agi- 
tate the  flask  to  thoroughly  mix  the  contents.  Place 


SILICON  10 1 

the  flask  in  the  position  shown,  letting  the  delivery 
tube  extend  to  the  bottom  of  the  cylinder.  Pour  mer- 
cury into  the  latter  until  the  end  of  delivery  tube  is  sub- 
merged, and  then  nearly  fill  the  cylinder  with  water. 

Gently  heat  the  flask  and  jo;bj<erv,er  the  gas  is  it  bubbles 
through  the  water  in  the  cylinder., ,  What  is  ,the  gelati- 
nous substance  formed?,'  ,'B&ve  iy6u  ,(hfer 'sVeft,' -the  com- 
pound before  ?  Where  is  the  fluosilicic  acid  ? 

After  the  experiment  has  been  running  for  some  time 
and  the  cylinder  of  water  contains  considerable  of  the 
gelatinous  precipitate,  disconnect  the  apparatus  and  fil- 
ter the  liquid  in  the  cylinder. 

241.  Test  the  action  of  the  clear  filtrate  from  Exp.  240 
with  litmus  paper.    What  sort  of  a  compound  is  in 
solution  ? 

To  a  small  portion  of  the  filtrate  add  BaCl2  solution. 
To  the  remainder  of  the  filtrate  add  KC1  or  KNO3  solu- 
tion and  allow  the  mixture  to  stand  for  30  minutes. 
Filter.  Dry  the  precipitate  between  pieces  of  filter  paper. 
Examine  carefully.  Test  a  small  portion  as  directed  in 
Exp.  235  and  another  portion  as  directed  in  Exp.  236. 

DETERMINATION  OF  THE  SPECIFIC  GRAVITY  OF  SAND. 

242.  Use  a  flask  with  a  mark  on  the  neck  as  in  Exp.  2  20. 
Weigh  the  flask  empty  (£).     Partially  fill  with  clean  dry 
sand  and  again  weigh  (S).    Now  add  water  until  the 
flask  is  filled  to  the  mark  on  the  neck;  again  weigh  (X). 

Empty  the  flask,  rinse  and  completely  fill  with  water 
alone  to  the  mark  on  the  neck.  Weigh  (W)> 

The  weight  of  substance  taken  equals  S  —  E.  The 
weight  of  water  needed  to  fill  the  flask  containing  the 
substance  equals  X  —  S.  The  weight  of  water  needed 


102  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

to  completely  fill  the  flask  equals  W  -  E.  The  differ- 
ence between  the  weight  of  water  in  the  two  cases  is 
(W  —  E)  —  (X  —  5)  and  equals  the  weight  of  a  volume 
of  water  equal  to  the  volume  of  the  substance  used. 
Therefore,  •  /  ;  :/•  j? 


BORON  (B;  n). 

243.  In  a  beaker  dissolve  10  gms.  of  borax  (Na2B4O7) 
in  40  cc.  of  boiling  water  and  to  the  solution  add  6  cc.  of 
concentrated  HCL     Allow  the  solution  to  cool.     What 
compound  crystallizes  out? 

Filter  and  wash  the  crystalline  mass  with  a  few  cubic 
centimeters  of  distilled  water.  Dry  by  pressing  be- 
tween pieces  of  filter  paper. 

Is  boric  acid  soluble  in  water?  If  so,  why  does  it 
crystallize  out  from  the  solution  prepared  above  ? 

244.  Make  a  solution  of  boric  acid  in  water.     Dip  a 
piece  of  turmeric  paper  into  the  solution.     Does  the 
paper  change  in  color?    Dry  the  paper  by  holding  for  a 
few  moments  high  above  a  burner  flame  or  by  putting  it 
around  the  neck  of  a  flask  in  which  water  is  being  boiled. 
What  is  the  color  of  the  paper  upon  drying?     What 
effect  does  a  drop  of  NH4OH  produce  on  the  color  ? 

245.  To  a  few  crystals  of  boric  acid  in  an  evaporating 
dish  add  a  few  cubic  centimeters  of  alcohol.     Ignite  the 
alcohol  and  observe  the  color  of  the  flame  until  the 
alcohol  is  entirely  burned.     Is  any  characteristic  color 
imparted  to  the  flame  by  the  boric  acid?     Explain. 

246.  In  separate  test  tubes,  add  a  solution  of  Na2B407 


BORON  103 

to  solutions  of  MnSO4,  CuS04,  and  CaCl2.     What  is  the 
nature  of  the  precipitates  ? 

247.  Make  a  borax  bead  (see  Exp.  162)  and  hold  in 
the  flame  until  perfectly  clear.     Allow  to  cool;    then 
separate  the  bead  from  the  wire  and  introduce  it  into  a 
test  tube  half  full  of  water.     Does  the  bead  dissolve? 
How  does  the  composition  of  the  borax  bead  differ  from 
that  of  ordinary  borax  ? 

248.  Repeat  Exps.  244  and  245,  using  borax  instead 
of  boric  acid,  and  adding  in  each  case  a  few  drops  of 
concentrated  H2SO4.     How  do  the  results  compare  with 
those  obtained  in  Exps.  244  and  245  ?     Why  is  it  neces- 
sary to  add  H2SO4  in  making  the  tests  with  borates, 
whereas  in  the  experiments  with  boric  acid  H2S04  was 
not  necessary  ? 

Problems,  (a)  To  make  15  kilos  of  a  21%  solution  of  sodium 
silicate,  what  weight  of  98%  pure  sand  and  what  weight  of  pure 
dry  sodium  carbonate  will  be  necessary? 

(b)  If  the  efficiency  of  the  electric  furnace  process  for  the  manu- 
facture of  carborundum  is  83%,  what  will  be  the  weight  of  the 
charge  necessary  to  produce  200  kilos  of  carborundum? 

(c)  If  75  liters  of  silicon  tetrafluoride  at  33°  and  740  mm.  pres- 
sure are  passed  through  water,  what  weight  of  dried  silica  and  how 
much  1 8%  fluosilicic  acid  will  be  obtained? 

(d)  From  2  tons  of  crystallized  borax  what  weight  of  pure 
dry  boric  acid  can  be  produced  ? 

(e)  What  volume  of  boron  trichloride  can  theoretically  be  ob- 
tained from  250  gms.  of  pure  boron  trioxide? 


CHAPTER  XII. 

PHOSPHORUS,  ARSENIC,  ANTIMONY  AND 
BISMUTH. 

PHOSPHORUS  (P;  31). 

249.  Dry  a  small  piece  of  phosphorus  by  means  of 
filter  paper.     Using   the  pincers,   place   the  piece   of 
phosphorus  on  a  dry  iron  dish  or  piece  of  asbestos  board 
and  allow  to  stand  in  the  air  until  it  ignites. 

250.  By  means  of  a  deflagrating  spoon,  burn  a  small 
piece  of  phosphorus  in  a  wide-mouth  bottle.     When 
combustion  is  complete,  withdraw  the  spoon  and  cover 
the  bottle  to  prevent  the  escape  of  the  white  fumes. 
Add  about  20  cc.  of  water  to  the  contents  of  the  bottle 
and  shake.     Test  the  reaction  of  the  water  towards  lit- 
mus paper.      (Reserve  the  solution  for  use  in  another 
experiment.) 

251.  Dissolve  a  piece  of  phosphorus  about  the  size  of 
a  grain  of  wheat  in  a  few  cubic  centimeters  of  C§2. 
Pour  the  solution  on  a  piece  of  filter  paper  and  allow  the 
C$2  to  evaporate.     (Do  not  get  the  solution  on  the  hands 
or  clothing.)     Why  does  the  phosphorus  take  fire  so 
readily  when  the  CS2  has  evaporated  ? 

252.  To  a  small  piece  of  phosphorus  in  a  test  tube  add 
3  or  4  cc.  of  concentrated  HN03  and  boil.     What  be- 
comes of  the  phosphorus?     (Save  the  solution  for  use 
in  another  experiment.) 

253.  Red  Phosphorus.     Carefully  dry  a  piece  of  phos- 
phorus about  half  the  size  of  a  pea  and  introduce  into  a 

104 


PHOSPHORUS  105 

test  tube.  Stopper  the  test  tube  tightly  with  an  ordi- 
nary cork.  Using  a  test  tube  holder,  hold  the  tube 
above  a  burner  with  the  flame  turned  low,  at  such  a 
distance  that  the  phosphorus  boils  slightly.  (HOOD.) 
Continue  to  heat  in  this  way  until  there  is  a  decided 
change  in  the  appearance  of  the  phosphorus.  (CAU- 
TION.) 

Allow  the  tube  to  cool  thoroughly.  Then  test  the 
solubility  of  the  product  in  CS2. 

Mention  several  ways  in  which  red  phosphorus  differs 
from  yellow  phosphorus. 

254.  Heat  a  bit  of  red  phosphorus  in  an  evaporating 
dish.     What  happens  ?    What  compound  is  formed  ? 

255.  Phosphorus  and  the  Halogens.     Under  the  hood, 
place  a  crystal  of  iodine  on  a  small  dry  piece  of  yellow 
phosphorus.     Allow  to  stand  for  a  moment.     Does  any 
action  take  place  ?    What  compound  is  formed  ? 

256.  Place  a  little  red  phosphorus  in  a  dry  test  tube 
standing  in  a  test-tube  rack  or  in  a  bottle  under  the  hood. 
(Do  not  hold  the  test  tube  in  the  hand.}     Into  a  second  test 
tube  pour  a  little  bromine.     Now  quickly  pour  the  bro- 
mine into  the  test  tube  containing  the  red  phosphorus. 
(CAUTION.) 

What  can  you  say  as  to  the  affinity  of  phosphorus  for 
the  halogens  ?  Make  a  list  of  names  and  formulas  of  all 
the  halogen  compounds  of  phosphorus  which  have  been 
studied  or  formed  in  experiments  up  to  the  present  time 
and  give  the  number  of  the  experiment  involved  in  each 
case. 

257.  Phosphine   (PH3).    Acidify  a  beaker  of  water 
with  HC1.     Place  under  the  hood  and  then  introduce 
a  piece  of  calcium  phosphide  (Ca3P2).    Allow  to  stand 


io6 


EXPERIMENTS   IN   GENERAL   CHEMISTRY 


for  several  minutes.     Describe  all  results  and  write  all 
equations  involved. 

258.  Under  the  hood  arrange  an  apparatus  as  shown 
in  Fig.  32.  Introduce  into  the  flask  25  cc.  of  strong 
NaOH  solution  and  five  or  six  small  pieces  of  yellow 
phosphorus.  Close  all  joints  tightly.  (Do  not  proceed 


FIG.  32. 

further  until  the  apparatus  has  been  approved  by  the  in- 
structor.) Connect  tube  a  with  a  gas  supply  and  pass 
gas  through  the  apparatus  to  displace  all  air.  Then 
shut  off  the  gas  and  apply  heat  carefully  to  the  flask. 
(Keep  the  hands  away  from  the  flask  after  heat  is  applied.) 
Phosphine  comes  from  the  exit  tube  and  burns  as  it 
strikes  the  air.  Note  the  white  rings  formed  when  the 
phosphine  burns.  What  is  the  composition  of  the  white 
fumes  ? 


PHOSPHORUS  107 

Continue  the  experiment  until  PHs  no  longer  comes 
from  the  apparatus.  Then  discontinue  the  heat  and 
immediately  turn  on  the  stream  of  gas  through  a  to  drive 
all  phosphine  from  the  apparatus.  After  this  is  done, 
cool  the  flask  by  introducing  100  cc.  of  water  into  it. 
(Pour  the  contents  of  the  flask  into  the  large  bottle 
labeled  " Hypophosphite  Solution.") 

(CAUTION  !  This  is  a  dangerous  experiment  and  should 
be  performed  with  the  greatest  care.) 

Is  there  more  than  one  compound  of  phosphorus  and 
hydrogen?  Describe  them  all. 

259.  Oxides  of  Phosphorus.     How  many  oxides  of 
phosphorus  are  there?     Which  of  the  oxides  have  you 
already  made?     Expose  a  small  amount  of  phosphorus 
pentoxide  (P205)  to  the  air.     What  happens?     Drop  a 
small  amount  of  the  dry  oxide  into  water.     What  do 
you  notice? 

From  these  two  tests  what  do  you  conclude  as  to  the 
affinity  of  PzO*,  for  water  ? 

Acids  of  Phosphorus. 

260.  Ortho-phosphoric  Acid  (H3PO4).     In  a  test  tube 
try  the  action  of  AgNOs  solution  on  a  solution  of  ordinary 
sodium  phosphate  (di-sodium  phosphate,  Na2HPO4). 

261.  To  about  i  cc.  of  Na^HPCX  solution  in  a  test 
tube  add  about  an  equal  volume  of  concentrated  HNO3 
and    then    a    like    volume    of    ammonium    molybdate 
((NH4)2Mo04)  solution.      If  no  precipitate  is  formed, 
warm  gently  for  a  time. 

Repeat  the  test,  using  the  solution  formed  in  Exp.  252 
instead  of  the  Na2HPO4  solution. 


108  EXPERIMENTS  IN  GENERAL  CHEMISTRY 

The  yellow  precipitate  formed  is  "  ammonium  phospho- 
molybdate."  This  is  the  best  test  for  phosphoric  acid. 

262.  To  a  few  cubic  centimeters  of  a  solution  of 
Na2HPO4  in  a  test  tube  add  solutions  of  NH4C1,  NH4OH 
and  MgS04.     What  is  the  compound  formed?    Why  is 
this  compound  of  considerable  importance  in  analytical 
chemistry  ? 

263.  When  the  yellow  precipitate  formed  in  Exp.  261 
has  settled,  carefully  decant  as  much  as  possible  of  the 
supernatant  liquid.     Then  add  strong  NH4OH  to  dis- 
solve the  yellow  residue.     To  the  solution  thus  formed 
add   solutions   of   NH4C1   and   MgSO4.     Compare   the 
precipitate  with  that  formed  in  Exp.  262. 

264.  Try  the  action  of  a  solution  of  sodium  phosphate 
on  solutions  of  copper  and  calcium. 

265.  Pyrophosphoric  Acid  (H4P2O7).     Prepare  sodium" 
pyrophosphate   (Na4P2O7)   by  strongly  heating   a  few 
crystals  of  di-sodium  phosphate  (Na2HPO4)  on  a  piece 
of  platinum  foil.     Dissolve  the  fused  salt  in  a  little  cold 
water. 

Test  a  portion  of  the  solution  with  AgNOs  solution. 

266.  Test  a  few  cubic  centimeters  of  a  pyrophosphate 
solution    with    HN03    and    ammonium    molybdate    as 
directed  in  Exp.  261.     Is  a  precipitate  formed?    Why? 

Repeat  the  test,  boiling  the  solution  with  the  HNO3 
for  a  moment  before  adding  the  ammonium  molybdate. 
Does  a  precipitate  form  now?  Why  is  it  necessary  to 
boil  the  solution? 

How  can  pyrophosphoric  acid  be  formed  directly  from 
orthophosphoric  acid? 

267.  Metaphosphoric  Acid  (HPO3).     Fuse  a  few  crys- 
tals of  microcosmic  salt  (sodium  ammonium  hydrogen 


PHOSPHORUS  IOQ 

phosphate,  NaNH4HP04)  on  a  piece  of  platinum  foil  or 
in  a  loop  at  the  end  of  a  platinum  wire.  Notice  the  odor 
of  the  gas  which  is  given  off.  Heat  until  effervescence 
ceases.  What  is  the  composition  of  the  remaining  salt? 
What  is  the  name  of  the  compound  ? 

268.  Dissolve  the  fused  salt  prepared  in  Exp.  267  in 
a  little  cold  water  and  test  a  small  portion  with  a  solution 
of  AgNO3.     To  a  second  portion  add  a  few  drops  of 
acetic  acid  and  then  a  little  egg  albumen  solution.     Why 
is  the  acid  necessary  ? 

Try  these  two  tests  on  the  solution  formed  in  Exp.  250. 
Which  of  the  phosphoric  acids  is  formed  when  P2O5  is  dis- 
solved in  cold  water?  .  How  can  metaphosphoric  acid 
be  formed  directly  from  orthophosphoric  acid  or  from 
pyrophosphoric  acid  ? 

269.  Phosphorous  Acid  (H3PO3).     Under  the  hood, 
pour  a  few  cubic  centimeters  of  phosphorus  trichloride 
(PCls)  into  a  flask  containing  30  cc.  of  water.     Agitate 
the  flask  to  hasten  reaction.     Notice  the  fumes  which 
are  evolved.     Test  their  action  towards  litmus  paper. 
What  are  the  fumes? 

Transfer  the  solution  in  the  flask  to  an  evaporating 
dish  and  evaporate  to  half  the  volume.  Why?  Allow 
the  resulting  syrupy  liquid  to  cool;  then  dilute  with 
about  10  cc.  of  distilled  water. 

Test  a  small  portion  of  the  solution  thus  formed  with 
AgNOs  solution.  Warm  gently  and  notice  the  precipi- 
tate. How  does  the  AgN03  test  for  H3PO3  differ  from 
that  with  H3PO4,  H4P2O7  and  HP03? 

270.  Place  the  remainder  of  the  phosphorous  acid  solu- 
tion in  an  evaporating  dish  and  evaporate  to  dryness. 
Note  all  phenomena  and  explain  fully. 


110  EXPERIMENTS   IN    GENERAL   CHEMISTRY 

271.  Hypophosphorous  Acid  (H3PO2).    Place  three  or 
four  pieces  of  yellow  phosphorus  in  25  cc.  of  baryta  solu- 
tion in  an  evaporating  dish  under  the  hood.     Place  the 
dish  on  a  ring  stand  and  apply  heat  to  bring  about 
reaction  (see  Exp.  258).     When  phosphine  is  no  longer 
evolved,  cool  and  filter  the  solution.     (//  is  well  to  burn 
the  filter  paper  and  residue,  as  it  may  contain  small  pieces 
of  phosphorus.) 

Pass  C02  through  the  clean  filtrate  to  precipitate  the 
excess  of  barium  in  the  solution.  Filter  and  evapo- 
rate the  filtrate  to  syrupy  consistency.  Upon  cooling, 
barium  hypophosphite  (Ba(H2P02)2)  should  crystallize 
out. 

272.  Dissolve  the  crystals  in  a  little  distilled  water, 
acidify  a  portion  with  HC2H3O2  and  add  AgNO3  solution. 
Watch  carefully  while  performing  this  test. 

Summary.  By  what  test  can  orthophosphoric  acid 
and  the  orthophosphates  be  distinguished  from  the  pyro 
and  meta  acids  and  their  salts?  By  what  tests  can 
pyrophosphoric  acid  and  its  compounds  be  distinguished 
from  the  ortho  and  meta  acids  and  their  salts?  What 
tests  can  be  used  to  distinguish  meta  phosphoric  acid 
and  its  salts  from  the  ortho  and  pyro  acids  and  their 
compounds  ? 

Show  by  equations  how  ortho,  pyro  and  meta  phos- 
phoric acids  can  be  made  from  P2O6  and  water.  Which 
of  these  acids  is  actually  made  when  P2C>5  is  dissolved 
in  cold  water  ? 

Show,  by  equations,  all  changes  which  take  place  when 
H3PO4  is  heated.  Which  acid  of  phosphorus  is  the  most 
stable  towards  heat?  Which  two  acids  of  phosphorus 
are  the  least  stable  when  heated  ? 


ARSENIC  III 

Problems,  (a)  What  volume  of  a  65%  solution  of  phosphoric 
acid  can  be  obtained  as  a  by-product  in  the  manufacture  of 
60  kilos  of  potassium  hypophosphite  ? 

(b)  From  two  liters  of  85%  phosphoric  acid,  what  weight  of 
metaphosphoric  acid  can  be  obtained? 

(c)  How  much  pure  yellow  phosphorus  will  be  necessary  in  the 
preparation  of  2  liters  of  phosphorus  trichloride  ? 

(d)  If  200  liters  of  phosphine,  at  180°  and  760  mm.  pressure, 
are  burned  in  an  inclosed  chamber  and  the  fumes  treated  with 
water,  what  volume  of  a  12%  solution  of  phosphoric  acid  can  be 
obtained  ? 

(e)  From  200  Ibs.  of  bone  ash,  running  94%  tricalcium  phos- 
phate, what  weight  of  phosphorus  can  be  produced  ? 

ARSENIC  (As;  75). 

273.  In  a  hard  glass  test  tube  strongly  ignite  a  small 
piece  of  arsenic.     Likewise  heat  a  little  of  a  mixture  of 
arsenic  trioxide  (As2O3)  and  powdered  charcoal.     Notice 
the  sublimate  in  each  case.     What  is  the  composition 
of  the  sublimate? 

274.  Heat  a  small  piece  of  arsenic  on  a  piece  of  char- 
coal with   the   oxidizing   blowpipe   flame.     Notice   the 
white  ring  which  forms  around  the  arsenic  and  at  some 
distance  from  the  latter. 

275.  Treat  a  little  powdered   arsenic   with   concen- 
trated HNO3  and  heat  to  boiling.     What  is  formed  ? 

Treat  a  little  powdered  arsenic  with  aqua  regia.     Does 
the  arsenic  dissolve  ? 
Try  the  solubility  of  arsenic  in  concentrated  HC1. 

276.  Test  the  solubility  of  small  amounts  of  arsenic 
trioxide  (A^Oa)  in  water,  in  concentrated  HC1,  in  HNOs 
and  in  a  solution  of  NaOH. 

277.  Pass  H2S  through  an  aqueous  solution  of 
Observe  carefully. 


112  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

Pass  H2S  through  an  aqueous  solution  of  As203  after 
first  adding  a  few  cubic  centimeters  of  HC1.  Why  is 
there  a  difference  in  these  two  tests?  Allow  the  pre- 
cipitate of  As2S3  to  settle;  then  decant  as  much  as 
possible  of  the  supernatant  liquid.  Treat  the  precipi- 
tate with  a  few  cubic  centimeters  of  (NH4)2S  solution. 
To  the  solution  thus  formed  add  concentrated  HC1  until 
acid  to  litmus  paper.  Note  all  changes  and  write  all 
equations. 

278.  In  separate  test  tubes  try  the  action  of  (i)  AgN03 
solution  and  (2)  CuSO4  solution  on  a  solution  of  an 
arsenite.     Note  the  color  of  each  of  the  precipitates. 

Likewise  try  the  action  of  an  arsenate  solution  with 
solutions  of  AgNOs  and  CuSO4.  Does  an  arsenate  give 
the  same  colors  as  an  arsenite  ? 

279.  Dissolve  a  small  amount  of  As203  in  NaOH, 
neutralize  a  small  portion  with  acetic  acid  and  test  with 
AgNOs  solution.     To  the  remainder  of  the  solution  add 
concentrated  HN03  until  strongly  acid  and  then  heat 
to  boiling.      Cool  under  the  faucet.     Add  NaOH  until 
alkaline  and  then  acetic  acid  until  slightly  acid.      Test 
a  portion  with  AgN03  solution.     What  compound  have 
you  made  in  this  experiment?     Write  all  equations. 

280.  Try  the  action  of  magnesia  mixture  (NH4OH 
+  NH4C1  +  MgS04)  on  a  solution  of  an  arsenate.    What 
other  acid  group  gives  a  similar  test  with  magnesia  mix- 
ture? 

Try  the  action  of  HN03  and  ammonium  molybdate 
solution  on  a  solution  of  an  arsenate.  Is  a  precipitate 
formed?  Heat  to  boiling  and  then  allow  to  stand 
quietly  for  a  few  moments.  How  does  this  test  differ 
from  that  with  a  phosphate  ? 


ARSENIC 


281.  Try  the  action  of  H2S  on  a  solution  of  an  arsenite 
which  has  been  acidified  with  HC1. 

Likewise  try  the  action  of  H2S  on  an  arsenate  solution 
which  has  been  acidified  with  HC1.  Is  there  a  difference 
in  these  two  tests  ?  Heat  the  arsenate  solution  to  boiling 
and  continue  to  pass  a  stream  of  H2S  through  the  solu- 
tion for  several  minutes.  Explain  fully  and  write  all 
equations. 


FIG.  33. 

282.  Marsh  Test  for  Arsenic.  Arrange  an  apparatus 
consisting  of  a  hydrogen  generator,  drying  tube,  filled 
with  dry  pieces  of  CaCl2,  hard  glass  tube  and  exit  tube 
drawn  to  a  fine  point,  as  shown  in  Fig.  33. 

Generate  hydrogen  by  the  action  of  HC1  or  dilute 
H2SC>4  on  pure  zinc.  Test  the  gas  coming  from  the  exit 
tube  and  as  soon  as  all  air  has  been  driven  from  the 


114  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

apparatus,  light  the  jet  after  first  wrapping  a  towel 
about  the  generator. 

Through  the  thistle  tube  now  add  a  few  cubic  centi- 
meters of  the  solution  to  be  tested  for  arsenic  (use  any 
arsenic  solution  in  this  experiment)  and  notice  the  change 
which  almost  immediately  is  produced  in  the  color  of 
the  flame. 

Collect  several  arsenic  spots  by  holding  pieces  of  cold 
porcelain  for  a  moment  in  the  arsenic  flame.  Save  these 
for  future  tests. 

Now  strongly  ignite  the  hard  glass  tube  through  which 
the  gases  pass  and  notice  the  black  deposit  of  arsenic 
which  is  produced  by  the  decomposition  of  the  arsine. 
Also  notice  the  change  in  color  of  the  flame  burning  at 
the  jet. 

Test  the  solubility  of  the  arsenic  spots  produced  above 
in  (i)  a  freshly  prepared  solution  of  sodium  hypochlo- 
rite,  and  (2)  concentrated  HNpa.  Record  all  observa- 
tions. 

(The  Marsh  Test  for  Antimony  (Exp.  287)  should 
preferably  be  performed  immediately  after  that  for 
arsenic.  If  this  is  not  done,  several  arsenic  spots  should 
be  reserved  for  comparison  with  the  antimony  spots.) 

ANTIMONY  (Sb;  85). 

283.  Strongly  ignite  a  small  piece  of  antimony  in  a 
hard  glass  test  tube.     Is  a  sublimate  formed  ? 

Melt  a  little  antimony  before  the  blowpipe  on  a  piece 
of  charcoal  and  drop  the  molten  globule  upon  a  piece  of 
manilla  paper  spread  out  on  the  desk. 

284.  In  a  test  tube  treat  a  little  powdered  antimony 


ANTIMONY  115 

with   concentrated  HN03.     Compare  with   the  corre- 
sponding experiment  with  arsenic. 

285.  Pour  a  few  cubic  centimeters  of  antimony  tri- 
chloride (SbCla)  solution  into  a  beaker  of  water.    What 
is  the  precipitate  formed?    Divide  into  two  portions 
and  treat  one  with  concentrated  HCL    Treat  the  other 
portion  with  a  saturated  solution  of  NaCl.     Can  you 
explain  these  tests? 

286.  Dilute   a   little   SbCla   solution   with  an   equal 
volume  of  water  and  pass  H2S  through  the  mixture. 
Test  the  solubility  of  a  portion  of  the  precipitate  in  con- 
centrated HC1.    Test  a  second  portion  with  a  solution 
of  (NH4)2S  and  warm  if  necessary.     To  the  solution  in 
(NH4)2S  now  add  concentrated  HC1  until  the  solution  is 
barely  acid.     Compare  the  precipitate  of  sulphide  thus 
formed  with  the  precipitate  first  obtained  with  SbCls. 
Are  they  the  same  ? 

Do  the  sulphides  of  arsenic  and  antimony  behave  alike 
when  treated  with  ammonium  sulphide? 

287.  Marsh  Test  for  Antimony.    Perform  the  Marsh 
test  for  antimony  in  exactly  the  same  manner  as  for 
arsenic.     Compare    the    arsenic    and    antimony    spots. 
Compare  the  deposits  produced  when  the  hard  glass  tube 
is  heated.     Which  comes  nearer  the  flame  ?     Do  HN03 
and  NaCIO  affect  the  antimony  spots  the  same  as  the 
arsenic  spots  ?    Write  all  equations. 

BISMUTH  (Bi;  208).^ 

288.  Make  a  mixture  of  a  little  bismuth  trioxide 
(Bi2O3),  and  dry  Na2CO3  and  heat  on  a  piece  of  charcoal 
with  the  reducing  flame  of  the  blowpipe.     Note  the 


n6 


EXPERIMENTS  IN  GENERAL  CHEMISTRY 


globule  of  metal  thus  formed.  How  does  it  compare 
with  lead  ? 

289.  Test  the  solubility  of  small  particles  of  bismuth 
in  (i)  HC1,  (2)  aqua  regia,  and  (3)  concentrated  HNOs. 
Write  ah1  equations.     How  does  the  action  of  HNOs  on 
bismuth  compare  with  its  action  on  the  other  members 
of  this  group  ? 

290.  Carefully  weigh  out  4  gms.  of  bismuth,  2  gms.  of 
lead  and  2  gms.  of  tin.     Place  the  metals  together  in  a 
small  beaker,  cover  with  water  and  heat  until  the  water 
boils.     Do  the  metals  melt? 

Transfer  the  metals  to  an  iron  crucible  and  heat 
strongly  to  completely  fuse  the  mixture.  Allow  to  cool; 
then  place  the  alloy  in  a  beaker  of  water  as  before  and 
heat  to  boiling.  What  is  the  melting  point  of  the  alloy  ? 
What  is  the  melting  point  of  each  of  the  metals  ? 

291.  Pour   a   little   bismuth 
trichloride  (BiCls)  solution  into 
a  beaker  of  water.     Add  concen- 
trated HC1. 

292.  Add  NH4OH  to  a  solu- 
tion of  bismuth.     What  is  the 
composition  of  the  precipitate? 
Is  it  an  acid  or  a  base  ?     Com- 
pare with  the  phosphorus  com- 
pound   (?)    having    a    similar 
formula. 

293.  Dilute  a  little  Bi(NO3)3 
FlG>  34'                  or  BiCls  solution  with  an  equal 

volume  of  water  and  pass  H2S  through  the  mixture 
until  precipitation  is  complete.  Filter;  wash  the  pre- 
cipitate on  the  filter  by  means  of  the  wash  bottle  as  shown 


BISMUTH  117 

in  Fig.  34.  (The  wash  bottle  can  be  easily  constructed 
and  should  always  be  kept  filled  with  water  ready  for  use.) 

Test  the  solubility  of  the  precipitated  bismuth  sul- 
phide (Bi2S3)  in  (NH4)2S  and  in  (NH4)2S,.  Does  it 
dissolve  ?  How  does  this  test  compare  with  similar  tests 
on  sulphides  of  arsenic  and  antimony? 

294.  To  a  solution  of  some  bismuth  compound  add  a 
little  freshly  prepared  solution  of  sodium  stannite.  Why 
is  this  a  good  test  for  bismuth?  What  compound  is 
formed? 

(The  sodium  stannite  solution  is  prepared  by  adding 
NaOH  solution,  a  little  at  a  time,  to  a  solution  of  stannous 
chloride  (SnCl2),  until  the  precipitate  which  is  at  first 
formed  redissolves.) 

Summary.  Point  out  the  ways  in  which  bismuth, 
antimony  and  arsenic  appear  to  be  similar.  Is  bismuth 
ever  acid  in  its  chemical  behavior  ?  Is  arsenic  ever  basic 
in  its  behavior?  How  does  antimony  stand  with  regard 
to  these  two?  Make  a  table  showing  the  oxides,  chlo- 
rides and  hydrides  of  each  of  the  elements  of  this  group. 

Problems,  (a)  How  much  white  arsenic  can  be  produced  from 
500  Ibs.  of  realgar  ?  From  500  Ibs.  of  orpiment  ? 

(b)  What  volume  of  H2S  at  standard  conditions  will  be  neces- 
sary to  completely  precipitate  the  arsenic  from  200  gms.  of  a  15% 
solution  of  sodium  arsenate,  considering  that  only  20%  of  the 
H2S  is  lost? 

(c)  From  25  gms.  of  white  arsenic,  what  volume  of  arsine  at 
o°  and  760  mm.  can  theoretically  be  produced  ? 

(d)  What  weight  of  iron  will  be  required  to  completely  reduce 
12  tons  of  pure  stibnite? 

(e)  What  volume  of  air  at  20°  and  765  mm.  will  be  required  in 
roasting  500  kilos  of  orpiment  (containing  85%  As2S3)  to  the 
oxide? 


CHAPTER  XIII. 

THE  ALKALIES  AND  AMMONIUM. 

LITHIUM  (Li;  7). 

295.  Thoroughly  clean  the  end  of  a  piece  of  platinum 
wire  by  dipping  into  concentrated  HC1  and  then  heating 
in  the  hottest  part  of  the  Bunsen  flame.    When  clean, 
it  will  not  color  the  flame.     Then  dip  the  clean  wire  into 
a  solution  of  lithium  chloride  (LiCl)  and  again  hold  in 
the  flame.     What  is  the  color  of  the  lithium  flame  ? 

296.  Look    through    the    spectroscope    towards    the 
window  and  carefully  focus  the  instrument  so  that  a 
sharp  image  is  produced.     Then  look  through  the  spec- 
troscope at  the  lithium  flame.    How  many  lines  has  the 
spectrum  of  lithium  as  seen  through  the  small  spectro- 
scope?    Draw  a  diagram  of  the  spectrum  showing  the 
relative  position  of  the  lithium  line. 

SODIUM  (Na;  23). 

297.  Place  a  small  piece  of  metallic  sodium  on  a  watch 
glass  and  allow  to  stand  exposed  to  the  air  for  several 
days.     Observe  the  various  changes  which  take  place. 
When  crystals  are  at  length  formed,  try  the  action  of 
HC1  upon  them. 

Write  all  equations  and  explain  all  changes  which 
have  taken  place. 

298.  Drop  a  piece  of  metallic  sodium  about  half  the 
size  of  a  pea  into  a  beaker  containing  about  25  cc.  of 
water.     What  gas  is  liberated?    Why  does  the  sodium 

118 


SODIUM  119 

float  on  the  water?    Does  it  burn  when  it  reacts  with 
water  ?     (Reserve  the  solution  for  use  in  Exp.  300.) 

299.  Dissolve  about  10  gms.  of  crystallized  sodium 
carbonate  (Na2CO3)  in  50  cc.  of  hot  water.    To  a  small 
portion  of  this  solution  add  HC1.     Does  it  effervesce? 
Why? 

In  a  mortar  mix  about  10  gms.  of  Ca(OH)2  with  enough 
water  to  form  a  thin  paste  and  add  this  mixture  (milk  of 
lime)  to  the  solution  of  sodium  carbonate.  Filter  and 
evaporate  the  filtrate  to  about  one-third  its  volume. 
Filter  again  if  not  clear.  Test  a  small  portion  of  the 
solution  with  HC1.  Does  it  effervesce?  Why?  What 
compound  is  in  solution  ? 

300.  Compare  the  solutions  obtained  from  Exps.  298 
and  299.     Test  each  solution  with  red  and  blue  litmus 
paper.     Do  the  solutions  behave  alike?     Try  the  action 
of  a  portion  of  each  solution  on  a  solution  of  FeCl^. 
Drop  a  little  phenolphthalein  into  a  portion  of  each 
solution. 

301.  Mix  about  50  gms.  of  Na2SO4,  25  gms.  Ca(OH)2 
and  200  cc.  of  water.     Heat  to  boiling.     Filter  rapidly 
through  a  plaited  filter.    What  compound  is  in  solution 
in  the  filtrate  ?    What  is  the  residue  ? 

Divide  the  clear  nitrate  into  two  equal  portions. 
Saturate  one  portion  with  CO2,  filter  if  necessary,  and 
then  add  the  other  portion.  Filter  again  if  the  solution 
is  not  clear.  Evaporate  to  about  half  the  volume  and 
allow  to  stand  quietly  to  crystallize.  What  is  the  com- 
position of  the  crystals?  What  is  the  object  of  pro- 
ceeding in  the  above  manner  ? 

Dry  the  crystals  between  filter  papers.  Test  a  small 
portion  with  HC1.  What  does  this  prove  ? 


120 


EXPERIMENTS   IN   GENERAL  CHEMISTRY 


302.  Solvay  Soda  Process.  Arrange  an  apparatus  as 
shown  in  Fig.  35.  Generate  NH3  in  the  flask  on  the 
ring  stand  by  boiling  a  strong  solution  of  NH4OH. 
Pass  the  NH3  through  25  cc.  of  a  saturated  solution  of 


FIG.  35. 

NaCl  until  it  is  saturated  with  the  gas.  Now  remove 
the  NH3  generator  and  pass  CO2  from  the  other  generator 
through  the  solution  until  saturated  and  a  white  pre- 
cipitate is  formed.  What  is  the  nature  of  this  precipi- 
tate? Is  it  soluble  in  water?  Test  a  small  portion 
with  HC1.  How  can  this  compound  be  changed  into 
Na2C03? 

303.  Le  Blanc  Soda  Process.  Make  an  intimate  mix- 
ture of  six  parts  of  dry  Na2S04,  four  parts  of  powdered 
CaCO3  and  one  part  powdered  charcoal.  Grind  them 
together  in  a  mortar.  Fuse  a  portion  of  the  mixture  on 
a  piece  of  platinum  foil. 


SODIUM  121 

Allow  to  cool,  extract  with  a  little  hot  water,  and  fil- 
ter the  solution.  What  does  the  nitrate  contain  ?  Test 
a  portion  of  it  with  HC1.  To  another  portion  add  a  few 
drops  of  baryta  or  lime  water. 

304.  Insoluble  Sodium  Salt.    To  a  few  cubic  centi- 
meters   of    a    solution    of    potassium    pyroantimonate 
(K2H2Sb207)  add  a  few  cubic  centimeters  of  a  strong 
solution  of  NaCl.     Allow  the  tube  containing  the  mix- 
ture to  stand  for  some  time.     What  is  the  composition 
of  the  crystals  which  are  formed  and  which  cling  to  the 
test  tube  ?    How  many  insoluble  salts  has  sodium  ? 

305.  Heat  a  few  good-sized  crystals  of  NaCl  (rock 
salt)    in  a   test   tube.      Explain    the    phenomena   ob- 
served. 

In  a  porcelain  mortar  powder  a  large  crystal  of  NaCl. 
Place  the  fine  powder  on  a  piece  of  platinum  foil  and 
heat  strongly.  What  happens?  Are  there  any  little 
explosions  ? 

306.  Dip  a  clean  platinum  wire  into  a  solution  of  NaCl 
and  hold  in  the  flame.     What  color  is  imparted  to  the 
flame  by  sodium  and  its  compounds  ? 

Repeat  the  experiment  and  examine  the  sodium  flame 
as  it  appears  through  a  piece  of  blue  glass. 

307.  Examine  the  sodium  flame  through  the  spectro- 
scope and  make  a  diagram  of  the  spectrum  showing  the 
sodium  line.     Mix  a  little  sodium  solution  and  lithium 
solution  and  make  a  spectroscopic  test  of  the  mixture. 
Can  you  recognize  the  lines  of  both  elements  ? 


122  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

POTASSIUM  (K;  39). 

308.  Drop  a  small  piece  of  potassium  into  a  large 
beaker  of  water.     (CAUTION!     To  protect  the  eyes  cover 
the  beaker  with  a  piece  of  paper  or  cardboard.)     How  does 
this  reaction   compare  with  that  between  water  and 
sodium  ? 

Add  a  drop  of  phenolphthalein  to  the  solution.  What 
kind  of  a  compound  is  in  solution  ? 

309.  Dip  a  clean  platinum  wire  into  a  solution  of  KC1 
and  hold  in  the  Bunsen  flame.     What  color  do  potassium 
compounds  impart  to  the  flame  ? 

Make  a  mixture  of  NaCl  and  KC1  solutions  and  make 
a  flame  test  of  the  mixture.  Can  you  recognize  the 
potassium  flame?  Repeat  the  experiment,  looking  at 
the  flame  through  a  piece  of  blue  glass.  Look  at  a  plain 
potassium  flame  through  a  piece  of  blue  glass.  What 
flame  do  you  conclude  can  be  seen  through  the  blue 
glass  ? 

310.  Look  at  the  potassium  flame  through  the  spec- 
troscope and  draw  a  diagram  of  the  spectrum.     Make  a 
mixture  of  solutions  of  NaCl,  LiCl  and  KC1  and  examine 
the  spectrum  of  the  mixture.     Can  you  recognize  the 
lines  of  each  element? 

311.  Add  about  25  gms.  of  wood  ashes  to  about  50  cc. 
of  water  in  a  beaker.     Heat  gently  for   10  minutes. 
Filter  the  solution  and  evaporate  the  clear  nitrate  to 
dryness  in  a  porcelain  evaporating  dish. 

Test  the  residue  with  HC1.  What  happens  and  what 
does  it  signify?  Dip  a  clean  platinum  wire  into  the 
solution  in  HC1  and  hold  in  the  Bunsen  flame.  Also 
examine  the  flame  with  the  spectroscope.  What  com- 


POTASSIUM  123 

pound  was  extracted  from  the  wood  ashes?  What 
is  one  of  the  chief  compounds  in  the  ashes  of  sea 
plants  ? 

312.  Into  a  test  tube  containing  a  few  drops  of  bro- 
mine, drop  a  very  minute  piece  of  metallic  potassium. 
(CAUTION!  HOOD.)     Repeat,  using  sodium   instead   of 
potassium.     What  do  you  conclude  as  to  the  relative 
affinities  of  sodium  and  potassium  for  bromine  ? 

313.  In  separate  test  tubes  try  the  action  of  a  so- 
lution of  KOH  on  solutions  of  FeCls,  Pb(C2H302)2  and 
Cr2(S04)3.     Repeat,  using  NaOH  instead  of  KOH.    Do 
these  two  bases  act  alike  ? 

314.  Insoluble  Salts.    To  a  little  concentrated  KC1 
solution  in  a  test  tube  add  a  few  drops  of  platinum 
chloride  solution.     Allow  to  stand  a  few  moments.     Ex- 
amine the  precipitate.     Divide  the  mixture  into  two 
parts  and  test  one  with  hot  water.     Test  the  other  por- 
tion with  alcohol.     What  can  you  conclude  as  to  the 
solubility  of  potassium  platinum  chloride  (K2PtCle)  ? 

315.  Mix  about  equal  volumes  of  strong  KC1  solution 
and  tartaric  acid  solution.     Allow  the  mixture  to  stand 
undisturbed  for  10  minutes.     What  is  the  composition  of 
the  precipitate  which  is  formed  ?    Why  did  it  not  form 
immediately  ? 

Filter  and  wash  the  precipitate.  Fuse  a  part  or  all 
of  it  on  platinum  foil.  Cool  and  then  test  with  a  drop 
of  HCL  What  happens?  What  compound  has  been 
formed  ?  Can  you  write  the  equation  ? 

316.  Preparation    of    KNO3.     Dissolve    25    gms.    of 
crude  KC1  in  about  50  cc.  of  water  in  a  beaker.    Add 
the  calculated  weight  of  NaN03.     Filter  the  solution  if 
necessary  and  evaporate  the  nitrate  to  half  its  volume. 


124  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

Allow  to  stand  and  cool  quietly.  Filter;  then  dry  the 
crystals  between  pieces  of  filter  paper.  Dissolve  the 
crystals  in  about  15  cc.  of  boiling  water  and  allow  to 
cool  and  crystallize.  Dry  the  crystals. 

Examine  the  crystals  by  flame  test  and  by  means  of 
the  spectroscope.  Dissolve  a  trace  in  a  little  water  and 
test  for  nitric  acid  as  directed  in  Exp.  154,  page  71. 
What  is  the  most  probable  impurity  in  these  crystals? 
Did  this  impurity  appear  in  the  spectroscopic  test? 
Can  you  explain  why  two  soluble  salts  such  as  NaCl 
and  KNO3  can  be  separated  by  crystallization  ? 

317.  Oxidation  by  Means  of  KNO3.     Under  the  hood, 
heat  a  mixture  of  about  2  gms.  KNO3  and  i  gm.  powdered 
charcoal  in  an  iron  crucible.     Allow  to  cool,  extract  with 
water  and  filter.     Test  a  portion  of  the  solution  with  an 
acid.    What  is  proved  by  this  test  ? 

318.  Strongly  heat  about  2  gms.  KN03  and  i  gm.  of 
sulphur  in  an  iron  crucible.     (HOOD.)     Allow  to  cool, 
extract  with  water  and  test  a  portion   for  the  pres- 
ence of  sulphates.     Explain  the  action  and  write  equa- 
tions. 

Explain  by  equations  how  KNO3  oxidizes.  How  many 
available  oxygen  atoms  has  KNO3  ? 

319.  Potassium  Iodide    (KI).     Dissolve   25   gms.  of 
KOH  in  150  cc.  of  distilled  water  by  the  aid  of  heat. 
Add  iodine  in  very  small  quantities  at  a  time  and  with 
constant  stirring,  until  a  further  addition  causes  the 
liquid  to  remain  brown.     Concentrate  the  solution  by 
boiling,  add  about  50  gms.  of  powdered  charcoal  and 
transfer  to  an  evaporating  dish  or  large  crucible.    Evapo- 
rate to  dryness,  cover  and  ignite  for  20  minutes  at  a  dull 
red  heat.     Dissolve  the  mass  in  warm  water,  filter,  con- 


AMMONIUM  125 

centrate  and  set  aside  to  crystallize.     Purify  by  recrys- 
tallization  from  distilled  water. 

320.  Potassium  Chlorate  (KC1O3).     Slake  75  gms.  of 
lime,  mix  with  30  gms.  of  KC1  and  add  sufficient  water 
to  form  a  thin  paste.     Heat  almost  to  boiling  and  pass 
in  chlorine  until  no  more  is  absorbed  and  the  lime  has 
passed  into  solution.     Boil  for  an  hour,  passing  C02 
through  it  during  the  last  10  minutes,  and  filter  while 
hot. 

Evaporate  the  filtrate  to  100  cc.  and  set  aside  to 
crystallize.  Obtain  a  second  and  third  crop  of  crystals 
from  the  mother  liquor.  Purify  by  recrystallization. 
The  crystals  should  give  no  test  for  chlorides. 

AMMONIUM  (NH4). 

321.  To  a  little  sodium  amalgam  in  a  test  tube  add  a 
strong  solution  of  NH4C1.     Notice  the  odor  of  the  re- 
sulting compound.     Can  the  group  NH4  be  liberated? 
Why  has  this  group  been  given  a  name  ending  in  "urn  "  ? 

322.  Heat  a  small  amount  of  NH4C1  on  a  piece  of 
platinum  foil.     What  happens?     Try  (NH4)2SO4  in  the 
same  way.    Do  sodium  and  potassium  salts  behave  in  a 
similar  manner  when  heated?     (See  Exp.  305.) 

323.  Try  the  action  of  NaOH  on  a  solution  of  some 
ammonium  salt.     Note  the  odor  of  the  escaping  gas. 
Test  with  turmeric  paper.     Warm  the  mixture  if  neces- 
sary. 

324.  Make  a  flame  test  with  NH4C1.     What  color  is 
imparted  to  the  flame  by  ammonium  salts  ? 

325.  Try     the     action     of     ammonium     hydroxide 
(NH4OH)    on    solutions    of    FeCls,    Pb(C2H3O2)2,    and 


126  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

Cr2(S04)3.     Compare  results  with  those  obtained  from 
Exp.  313. 

326.  Test  for  an  Ammonium  Compound.     Place  a 
few  grams  of  Ca(OH)2  in  a  small  beaker  and  moisten 
with  the  solution  to  be  tested  for  ammonium  salts. 
Quickly  cover  the  beaker  with  a  watch  glass,  on  the  under 
side  of  which  is  a  piece  of  wet  turmeric  paper.     Allow  to 
stand  for  several  minutes. 

Try  this  test  on  a  number  of  ammonium  compounds. 

327.  Dissociation  of  NH4C1.     Place  a  loose  plug  of 
asbestos  fiber  in  the  middle  of  a  piece  of  large  glass  tubing. 
On  one  side  of  the  plug  and  close  to  it,  place  a  few  grams 
of  NH4C1.     At  each  end  of  the  tube  introduce  two  pieces 
of  moist  litmus  paper,  one  red  and  one  blue. 

Now  gently  heat  the  tube  directly  under  the  NH4C1 
to  volatilize  some  of  the  latter.  Carefully  observe  any 
changes  in  the  litmus  paper.  What  conclusion  can  you 
draw  from  the  results  of  this  experiment  ?  Who  was  the 
first  scientist  to  notice  this  phenomenon  ? 

Summary.  In  what  respects  does  ammonium  hydrox- 
ide differ  from  the  other  alkali  hydroxides?  In  what 
respects  is  ammonium  hydroxide  similar  to  the  other 
alkali  hydroxides?  Why  is  ammonium  grouped  with 
the  alkalies? 

How  can  solutions  of  (i)  lithium,  (2)  sodium,  (3)  po- 
tassium and  (4)  ammonium  be  distinguished  ?  How  can 
each  of  these  be  distinguished  in  presence  of  the  others  ? 

What  is  the  relation  between  " ammonia"  and  "am- 
monium compounds"  ?  Why  does  NH4OH  always  have 
the  odor  of  NH3? 

Problems,  (a)  To  prepare  10  tons  of  crystallized  sodium  car- 
bonate, how  much  sodium  chloride  is  necessary? 


AMMONIUM  127 

(b)  From  30  kilos  of  wood  ashes  containing  6%  of  K2CO3, 
what  weight  of  20%  KOH  solution  can  be  made  ? 

(c)  From  10  Ibs.  of  crude  ammonium  sulphate  (94%),  what  vol- 
ume of  dry  NH3  gas  can  be  prepared  at  15°  and  732  mm.  pressure? 

(d)  Considering  that  NH4C1  does  not  dissociate  when  vapor- 
ized, what  volume  will  260  gms.  of  NH4C1  occupy  when  com- 
pletely volatilized  at  600°  and  760  mm.  pressure  ? 


CHAPTER  XIV. 

THE  ALKALINE  EARTHS. 
CALCIUM  (Ca;  40). 

328.  Place  a  lump  of  lime  (CaO)  on  a  watch  glass  and 
add  water,  a  few  drops  at  a  time,  until  the  lime  slakes. 
The  water  can  be  very  conveniently  added  by  means  of 
a  wash  bottle  (Fig.  34).     Be  careful  to  avoid  an  excess 
of   water.     Is   heat   liberated   when    the   lime   slakes? 
Why? 

329.  Place  the  slaked  lime  prepared  above  in  a  5oo-cc. 
flask  and  add  about  350  cc.  of  distilled  water.     Cork 
tightly  and  allow  to  stand  with  occasional  shaking  for  an 
hour.     Filter  the  solution  into  a  clean  flask  and  stopper 
tightly.     Label  the  solution  "Lime  Water,"  and  reserve 
it  for  use  in  experiments  to  follow. 

Test  the  action  of  lime  water  towards  litmus  and 
turmeric  paper.  Test  a  few  cubic  centimeters  with  a 
drop  of  phenolphthalein  solution. 

330.  Place  a  small  piece  of  marble  in  a  porcelain 
crucible,  cover  and  ignite  over  the  blast  lamp  for  15 
minutes.     Allow  to  cool,  still  covered.     When  cool,  re- 
move the  cover  and  test  the  contents  of  the  crucible  for 
Ca(OH)2  by  means  of  wet  turmeric  paper  and  wet  red 
litmus  paper. 

Write  equations  to  show  what  happens  when  calcium 
oxalate  (CaC2O4)  is  ignited,  (ist)  gently,  and  (2nd)  with 
the  full  force  of  the  blast  lamp. 

331.  Treat  about  200  cc.  of  spent  liquid  from  a  C02 
generator  with  enough  slaked  lime  to  completely  neutral- 

128 


CALCIUM  129 

ize  the  acid  reaction.  Filter  the  solution  if  not  per- 
fectly clear.  To  the  filtrate  add  a  clear  solution  of 
Na2CO3  until  precipitation  is  complete.  Allow  the  pre- 
cipitate to  settle,  decant  the  supernatant  liquid  and  add 
distilled  water  to  the  white  precipitate.  Filter,  wash 
the  precipitate  on  the  filter,  and  allow  to  dry.  When 
dry  transfer  to  the  stock  bottle  labelled  "  Precipitated 
Calcium  Carbonate." 

Test  a  small  portion  of  the  precipitate  with  HC1. 
What  does  this  test  show? 

332.  Treat  a  few  cubic  centimeters  of  lime  water  with 
C02  until  the  precipitate  at  first  formed  redissolves. 
Explain  the  phenomenon  fully. 

Divide  the  solution  into  two  equal  portions.  Heat  one 
portion  to  boiling  for  a  moment.  To  the  other  portion 
add  clear  lime  water.  What  is  precipitated  in  each  of 
these  tests?  What  commercial  use  is  made  of  these 
reactions  ? 

333.  Test  separate  portions  of  any  calcium  solution 
with  solutions  of  ammonium  oxaiate  ((NH4)2C2O4)  and 
sodium  phosphate  (Na2HP04). 

334.  Make  a  flame  test  for  calcium,  using  a  solution 
of  CaCl2  or  any  other  calcium  solution  acidified  with 
HC1. 

Observe  the  spectrum  of  calcium  and  draw  a  diagram 
showing  the  various  lines. 

335.  To  50  cc.  of  a  clear  solution  of  CaCl2  add  dilute 
H2SC>4  until  precipitation  is  complete.     Filter  and  wash 
well  with  distilled  water.     Allow  to  drain.     Then  make 
a  hole  in  the  filter  paper  and  by  means  of  the  wash  bottle 
(Fig.  34)  wash  all  of  the  precipitate  into  a  clean  flask. 
Add  150  cc.  of  water,  cork  tightly  and  allow  to  stand 


130  EXPERIMENTS  IN  GENERAL  CHEMISTRY 

with  occasional  shaking.     Label  the  solution  "CaS04, 
Saturated  Solution,"  and  reserve  for  future  use. 

336.  To  a  few  cubic  centimeters  of  CaCl2  solution 
add  NH4OH  and  (NH4)2C03  solution  in  slight  excess. 
Filter.     To  the  clear  filtrate  add  a  few  drops  of  ammo- 
nium oxalate  ((NH4)2C204)  solution.     Does  a  precipitate 
form?     What  does  this  experiment  show  as  to  the  rela- 
tive solubility  of  the  carbonate  and  oxalate  of  calcium  ? 

337.  Heat   a   little   calcium   oxalate    (CaC204)    with 
concentrated  H2S04  in  a  test  tube  and  test  the  gas 
evolved  for  (i)  CO  and  (2)  C02.     Are  both  present? 
What  is  left  in  the  tube? 

338.  Plaster  of  Paris.     Powder  a  few  grams  of  gyp- 
sum and  heat  in  a  porcelain  dish,  stirring  the  powder 
with  a  thermometer  and  using  great  care  to  prevent 
the  temperature  rising  above  120°.     Why?    Allow  the 
mass  to  cool  thoroughly. 

Mix  the  cold  powder  with  enough  water  to  form  a 
thick  paste.  Lay  a  coin  on  a  glass  plate  and  pour  the 
paste  over  it.  Allow  to  stand  and  harden.  When 
perfectly  hard,  remove  the  coin  and  examine  the  im- 
pression. 

339.  Heat  a  second  and  smaller  portion  of  gypsum 
in  an  iron  crucible,  using  the  full  force  of  the  Bunsen 
flame.     Allow  to  cool ;  then  mix  with  water  and  allow  to 
stand.     Does  the  mass  set?     In  what  does  this  latter 
preparation  differ  from  plaster  of  Paris  ? 

Phosphate  Fertilizers. 

340.  Insoluble      Calcium     Phosphate      (Ca3(P04)2). 
Treat  a  little  powdered  apatite  or  bone  ash  with  a  few 
cubic  centimeters  of  distilled  water  and  heat  to  boiling. 


CALCIUM  131 

Filter  and  refilter  until  the  solution  is  perfectly  clear. 
Test  the  clear  filtrate  for  phosphates  by  means  of  HN03 
and  ammonium  molybdate  solution  (see  Exp.  261).  Is 
the  apatite  soluble  in  water  ? 

To  prove  that  the  mineral  contains  phosphoric  acid, 
dissolve  a  small  portion  in  concentrated  HNO3  and  again 
test  with  ammonium  molybdate  solution. 

341.  Preparation   of    Superphosphate    (Ca(H2PO4)2). 
Treat  20  gms.  of  powdered  apatite  or  bone  ash  with 
5  cc.  of  concentrated  H2SO4  in  a  porcelain  dish.     Warm 
gently  for  about  10  minutes. 

Add  a  small  portion  of  the  resulting  mass  to  15  or 
20  cc.  of  water  and  test  the  solution  thus  formed  for 
phosphates.  (Save  the  remainder  of  the  mass  for  use  in 
Exp.  342.) 

Has  the  action  of  the  H2SO4  caused  the  insoluble 
phosphate  to  become  soluble?  Show  by  an  equation 
how  this  was  done  ? 

342.  Reverted  Phosphate  (CaHPO4).    To  the  mass, 
containing  superphosphate  which  was  left  from  Exp.  341, 
add  25  cc.  of  water  and  about  30  gms.  of  Ca(OH)2.     Heat 
to  boiling,  allow  to  stand  15  minutes,  and  then  filter. 
(Save  both  the  nitrate  and  the  residue  on  the  filter.) 

Test  a  portion  of  the  filtrate  for  phosphates.  Do  you 
get  a  test?  Why?  What  has  become  of  the  super- 
phosphate? (Equation.)  Where  is  the  reverted  phos- 
phate ?  Is  reverted  phosphate  soluble  in  water  ? 

Transfer  the  filter  containing  the  residue  to  a  small 
flask  and  treat  with  20  cc.  of  a  solution  of  ammonium 
citrate.  Shake  to  thoroughly  mix  the  contents.  Warm 
gently  for  a  few  moments.  After  standing  for  15  min- 
utes, filter  and  test  the  clear  filtrate  with  HNO3  and 


132  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

ammonium  molybdate  solution.  What  is  present  ?  Can 
you  explain  everything  in  this  experiment?  Of  what 
use  is  reverted  phosphate  in  a  fertilizer  ? 

EXAMINATION  OF  FERTILIZERS. 

(Qualitative.) 

343.  Superphosphate.  Treat  about  10  gms.  of  the 
fertilizer  with  30  cc.  of  water  in  a  small  flask.  Allow  to 
stand,  with  occasional  shaking,  for  20  minutes.  Filter 
and  test  the  clear  nitrate  for  phosphates.  (Use  the 
residue  in  the  test  for  "reverted  phosphate.") 

Reverted  Phosphate.  Wash  the  residue  in  the  filter 
(from  the  preceding  test)  several  times  with  water  to  re- 
move the  last  traces  of  superphosphate.  Then  transfer 
the  filter  paper  and  residue  to  a  small  flask,  treat  with 
20  cc.  of  ammonium  citrate  solution,  shake,  and  gently 
warm  for  about  20  minutes.  Filter  and  test  the  clear 
filtrate  for  phosphates  (reverted  phosphate).  (Save  the 
residue  on  the  filter  for  the  next  test.) 

Insoluble  Phosphate.  Wash  the  residue  from  the  pre- 
ceding test  with  warm  water  to  remove  the  last  traces 
of  ammonium  citrate  solution  and  reverted  phosphate. 
Dissolve  a  portion  of  the  clean  residue  in  a  little  con- 
centrated HN03  and  test  for  phosphates  (insoluble  phos- 
phate) . 

Potash.  Treat  a  bit  of  the  fertilizer  on  a  watch  glass 
with  a  drop  or  two  of  concentrated  HC1.  By  means  of 
a  clean  platinum  wire,  make  a  flame  test  of  the  resulting 
solution  and  observe  the  flame  through  a  piece  of  blue, 
glass.  Can  you  detect  the  potassium  flame? 

Repeat  and  observe  the  flame  through   the  spectro- 


STRONTIUM  133 

scope.  Can  you  recognize  the  potassium  lines?  What 
other  lines  are  visible  ? 

In  a  small  beaker  treat  about  i  gram  of  the  fertilizer 
with  25  cc.  of  water  and  3  gms.  of  Ca(OH)2.  After 
standing  5  minutes,  filter.  To  the  filtrate  add  NH4OH 
in  excess  and  a  solution  of  (NH4)2C2O4  until  precipita- 
tion is  complete.  Heat  to  boiling  and  filter  while  hot. 

Evaporate  the  clear  filtrate  to  dryness  in  a  porcelain 
evaporating  dish  and  heat  gently  until  white  fumes 
(ammonium  salts)  no  longer  come  off.  Cool;  dissolve 
the  residue  in  a  few  drops  of  water,  filter  if  necessary 
through  a  very  small  wet  filter,  and  to  the  clear  filtrate 
add  a  few  drops  of  platinum  chloride  solution  (H2PtCl6) . 
(See  Exp.  314.)  A  yellowish  red  precipitate  is  K2PtCl6 
and  shows  presence  of  potash. 

Ammoniacal  Nitrogen.  This  means  nitrogen  which 
is  present  in  the  form  of  ammonium  salts.  Test  as 
described  in  Exp.  326,  page  126. 

Nitrogen  as  Nitrate.  The  above  test  does  not  show 
nitrogen  present  as  nitrate.  All  nitrates  are  soluble  in 
water.  Treat  a  small  portion  of  the  fertilizer  with  dis- 
tilled water,  filter  and  test  the  filtrate  for  nitrates  as 
directed  in  Exp.  154,  page  71. 

STRONTIUM  (Sr;  87). 

344.  In  separate  test  tubes  try  the  action  of  solutions 
of  the  following  substances  on  a  solution  of  strontium: 
ammonium    carbonate,    sodium    phosphate,   potassium 
chromate  and  dilute  H2SC>4. 

345.  Make  a  flame  test  for  strontium.    What  other 
elements  give  similar  flame  tests  ? 

Observe  the  spectrum  of  strontium  and  draw  a  diagram 


134  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

showing  the  more  important  lines.  Compare  the 
diagram  with  those  of  elements  which  give  a  flame  test 
similar  to  strontium. 

346.  To  about  15  cc.  of  a  strontium  solution  add 
dilute  H2S04  until  precipitation  appears  to  be  complete. 
Filter  and  wash  the  precipitate.     Make  a  hole  in  the 
paper  and  by  means  of  the  wash  bottle,  wash  the  pre- 
cipitate into  a  clean  flask.     Treat  with  about  50  cc.  of 
distilled  water.     Cork  tightly  and  allow  to  stand  with 
occasional    shaking.      Label   "  SrSO4,    Saturated  Solu- 
tion," and  reserve  for  future  use.     (Exp.  349.) 

BARIUM  (Ba;  137). 

347.  In  separate  test  tubes  try  the  action  of  solutions 
of  the  following  substances  on  a  solution  of  some  barium 
salt:    ammonium   oxalate,    sodium    phosphate,    sodium 
sulphate,   ammonium   carbonate,   potassium   chromate, 
and   dilute   H2S04.     Compare   the   results   with   those 
obtained  from  Exps.  333  and  344. 

348.  Make  a  flame  test  for  barium,  using  BaCl2  solu- 
tion and  a  few  drops  of  HC1. 

Observe  the  spectrum  of  barium  and  draw  a  diagram 
showing  the  lines. 

349.  Relative  Solubility  of  the  Alkaline  Earth  Sul- 
phates.    Filter  a  little  of  the  SrS04  solution  prepared 
in  Exp.  346.     To  a  few  cubic  centimeters  of  this  clear 
solution  add  an  equal  volume  of  some  other  strontium 
solution.     Is  there  a  precipitate  formed?    Why? 

Filter  a  little  of  the  CaSO4  solution  from  Exp.  335  and 
test  a  few  cubic  centimeters  of  this  with  an  equal  volume 
of  some  strontium  solution  (other  than  SrSO4).  Is  there 
a  precipitate  formed?  Why  is  a  precipitate  produced 


BARIUM  135 

by  a  saturated  solution  of  CaSCX  and  not  by  a  satu- 
rated solution  of  SrS04?  What  do  you  conclude  as  to 
the  relative  solubility  of  the  sulphates  of  calcium  and 
strontium  ? 

Add  a  few  cubic  centimeters  of  the  saturated  CaSC^ 
solution  to  a  solution  of  some  barium  salt.  Does  it 
produce  a  precipitate  ?  Try  the  action  of  a  little  of  the 
saturated  SrSCX  solution  on  a  barium  solution.  Does 
this  also  cause  a  precipitate?  What  can  you  conclude 
as  to  the  solubility  of  BaS04?  Wliich  of  the  alkaline 
earth  sulphates  is  the  least  soluble  and  which  is  the  most 
soluble?  Are  any  of  them  as  soluble  as  the  alkali 
sulphates  ? 

DETERMINATION  OF  THE  NUMBER  OF  MOLECULES  OF 
WATER  OF  CRYSTALLIZATION  IN  BARIUM  CHLORIDE. 

(Quantitative.) 

350.  The  determination  of  the  number  of  molecules 
of  water  of  crystallization  in  barium  chloride  is  carried 
out  in  exactly  the  same  manner  as  was  the  determina- 
tion of  the  number  of  molecules  of  water  of  crystalliza- 
tion in  gypsum.  (See  Exp.  44,  page  31.) 

Summary.  In  what  general  principles  do  the  alkaline 
earths  differ  from  the  alkalies?  What  can  you  say  as 
to  the  relative  solubility  of  the  compounds  of  these  two 
groups  of  metals  ? 

In  what  respects  does  lithium  somewhat  resemble  the 
alkaline  earths?  In  what  respects  does  it  more  closely 
resemble  the  alkalies?  What  one  point  is  alone  suffi- 
cient to  cause  lithium  to  be  placed  in  the  group  with  the 
alkalies  ? 

By  what  test  or  tests  can  each  of  the  alkaline  earths 


136  EXPERIMENTS   IN    GENERAL   CHEMISTRY 

be  distinguished?  Mention  tests  by  means  of  which 
each  of  the  alkaline  earths  can  be  detected  in  presence 
of  the  other  two. 

Problems,     (a)  How  many  kilos  of  lime  can  be  prepared  from 
4  tons  of  pure  CaCO3? 

(b)  What  volume  of  15%  HC1  would  be  required  to  dissolve 
1840  gms.  of  pure  CaC03? 

(c)  How  much  gypsum  is  necessary  for  the  preparation  of  500 
Ibs.  of  plaster  of  Paris  ? 

(d)  By  heating  15  liters  of  a  15%  solution  of  calcium  bicar- 
bonate, what  volume  of  CO2  (standard  conditions)  will  be  liberated  ? 

(e)  How  much  lime  will  be  necessary  to  completely  soften 
15  cubic  meters  of  water  in  which  the  hardness  is  due  entirely 
to  calcium  bicarbonate  and  in  which  this  compound  is  present  to 
the  extent  of  i.8%? 


CHAPTER  XV. 
MAGNESIUM,   ZINC,   CADMIUM   AND    MERCURY. 

MAGNESIUM  (Mg;  24). 

351.  Burn  a  small  piece  of  magnesium  ribbon  and 
allow  the  oxide  to  fall  upon  a  watch  glass.     Add  a  drop 
of  water,  allow  to  stand  for  a  moment  and  test  the 
reaction  towards  litmus  paper.     What  do  you  conclude 
from  this  test  ? 

352.  Test  the  solubility  of  magnesium  in  the  dilute 
acids.     Can  you  identify  the  gaseous  products  formed  ? 
Write  all  equations. 

353.  Test  separate  portions  of  a  solution  of  MgCl2 
or  MgSO4  with  solutions  of  the  following  substances: 
NH4OH,  NaOH,  and  Na2CO3. 

354.  Add  a  few  cubic  centimeters  of  NaOH  solution 
to  a  solution  of  magnesium  (MgCl2  or  MgSO4).     Is  there 
a  precipitate  formed?     Now  treat  the  mixture  with  a 
strong  solution  of  NH4C1.     What  happens  ?     Why  ? 

355.  Evaporate  25  cc.  of  MgCl2  solution  nearly  to 
dryness  in  a  porcelain  evaporating  dish.     Test  the  vapors 
occasionally  by  means  of  litmus  paper.     Do  they  have 
any   reaction?     Test   the   residue   in   the    evaporating 
dish  with  litmus  paper. 

356.  Mix  25  cc.  of  MgCl2  solution  with  an  equal 
volume  of  NH4C1  solution  in  an  evaporating  dish  and 
evaporate  nearly  to  dryness,  as  in  the  preceding  experi- 
ment.    Likewise  test  the  vapor  and  the  residue  with 

137 


138  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

litmus  paper.  Why  do  the  results  obtained  differ  from 
those  obtained  from  the  preceding  experiment  ?  Of  what 
commercial  importance  are  these  experiments  ? 

357.  Dissolve  5  gms.  of  MgO  in  dilute  H2SO4,  avoid- 
ing an  excess  of  acid.     Filter  the  solution  if  not  perfectly 
clear,  and  evaporate  to  a  small  volume.     Allow  to  stand 
quietly  and  cool.     Drain  the  crystals  on  a  filter,  redis- 
solve  them  in  a  little  hot  water,  and  allow  to  recrys- 
tallize. 

Examine  the  crystals.  Note  their  taste.  How  many 
molecules  of  water  of  crystallization  has  MgSO4?  To 
what  class  of  compounds  does  it  belong? 

358.  Make  a  mixture  of  equal  parts  of  solutions  of 
NH4C1,  NH4OH  and  some  magnesium  salt  (MgCl2  or 
MgSO4).     What  name  is  applied  to  this  mixture  ?    Why 
is  the  mixture  a  good  reagent  for  phosphates?    Try  the 
action  of  the  mixture  on  any  phosphate  solutions  you 
find  on   the .  reagent   shelves.     Is   a  white  precipitate 
formed  in  each  case? 

Try  the  action  of  an  acid  on  a  portion  of  the  white 
precipitate  of  ammonium  magnesium  phosphate  thus 
formed.  To  the  solution  in  acid  now  add  an  excess  of 
NH4OH. 

What  gas  would  be  liberated  if  you  heated  the  ammo- 
nium magnesium  phosphate?  What  residue  would  be 
left  ?  Write  the  equation  for  this  reaction. 

359.  Dissolve  a  piece  of  magnesium  ribbon  in  dilute 
HNO3  and  evaporate  the  solution  to  dryness  in  a  porce- 
lain dish.     (HOOD.)     Strongly  ignite  the  dish  and  con- 
tents for  a  few  moments.     When  cool,  treat  with  a  few 
drops  of  water  and  test  with  a  piece  of  red  litmus  paper. 
Compare  with  Exp.  351. 


ZINC  139 

ZINC  (Zn;  65). 

360.  Test  the  solubility  of  metallic  zinc  in  dilute  and 
concentrated  HC1,  in  dilute  and   concentrated  H2S04 
and  in  HN03.     Does  it  make  any  difference  whether  the 
acid  is  dilute  or  concentrated?    What  gaseous  products 
are  formed  in  each  case  ? 

361.  Heat  a  small  piece  of  zinc  on  charcoal  with  the 
oxidizing  flame  of  the  blowpipe.     Notice  the  deposit  of 
zinc  oxide   (ZnO)   formed  on  the  charcoal.     Note  its 
color  when  hot  and  when  cold.     Is  there  a  difference  ? 

Moisten  the  ZnO  on  the  charcoal  with  a  drop  of  a 
solution  of  cobalt  nitrate  (Co(N03)2)  and  again  heat 
with  the  blowpipe.  What  color  is  produced?  What 
is  the  composition  of  the  colored  compound  ? 

362.  To  a  solution  of  ZnSO4  gradually  add  NaOH 
solution  until  in  excess.     Note  all  the  changes  and  write 
all  equations.     Name  the  two  zinc  compounds  which 
have  been  formed  in  this  experiment. 

363.  To  a  portion  of  the  solution  formed  in  Exp.  362 
add  dilute  HC1  a  little  at  a  time  until  the  solution  is  acid 
to  litmus.     Explain  all  changes  and  give  names  of  com- 
pounds formed. 

364.  Treat     separate    portions    of    ZnS04    solution 
with  solutions   of   the   following   substances:    Na2CO3, 
K4Fe(CN)6,  and  Na2HP04. 

365.  Add  a  few  drops  of  HC1  to  a  test  tube  half  full 
of  ZnS04  solution  and  pass  H2S  through  the  mixture. 
Is  there  a  precipitate  formed  ? 

To  a  second  portion  of  ZnSO4  solution  add  (NH^S 
solution.  Divide  the  mixture  into  two  portions  and  to 
one  add  concentrated  HC1  and  to  the  other  acetic  acid 


140  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

(HC2H302).     What  can  you  say  as  to  the  solubility  of 
ZnS? 

366.  Heat  a  few  crystals  of  ZnSO4  in  a  dry  test  tube. 
Does   ZnSO4    contain   water   of   crystallization?    How 
many  molecules?     To  what  class  of  compounds  does  it 
belong  ? 

CADMIUM  (Cd;  112). 

367.  Heat  a  small  piece  of  cadmium  on  a  piece  of  char- 
coal with  the  oxidizing  flame  of  the  blowpipe.     What 
color  is  cadmium  oxide  (CdO)  ? 

368.  Try  the  action  of  solutions  of  the  following  sub- 
stances on  separate  portions  of  a  solution  of  cadmium 
sulphate  (CdS04):  Na2CO3,  K4Fe(CN)6  and  Na2HPO4. 
Compare  with  the  results  obtained  from  Exp.  364. 

369.  To  a  little  CdSO4  solution  add  NaOH  solution 
a  little  at  a  time  until  in  excess.     Does  the  cadmium 
solution  behave  like  a  zinc  solution  when  treated  in  this 
way?     (See  Exp.  362.) 

370.  Pass  H2S  through  a  solution  of  CdS04.     Try 
the  solubility  of  separate  portions  of  the  precipitate  in 
(i)  dilute  HC1,  (2)  concentrated  HC1  and  (3)  HC2H3O2. 
Compare  with  Exp.  365. 

How  does  the  solubility  of  CdS  compare  with  the 
solubility  of  ZnS? 

MERCURY  (Hg;  200). 

371.  In  a  hard  glass  test  tube  heat  a  little  of  any 
mercury    compound   with   about    twice   as   much   dry 
Na2CO3.     What  is  the  composition  of  the  sublimate? 
Rub  it  with  a  glass  rod. 

372.  Place  a  little  cinnabar  (HgS)  in  the  middle  of  a 
piece  of  hard  glass  tubing  open  at  both  ends,  and  clamp 


MERCURy  141 

the  tube  in  a  slightly  inclined  position.     Strongly  heat 
the  tube  at  the  point  just  below  the  HgS. 

Notice  the  sublimate  formed.  Also  notice  the  odor 
of  any  gases  coming  from  the  upper  end  of  the  tube. 
Test  their  action  on  wet  blue  litmus  paper. 

373.  Using   minute    globules    of    mercury,    test    the 
solubility  of  the  latter  in  both  dilute  and  concentrated 
HC1,   dilute  and  concentrated  H2S04,  and  dilute  and 
concentrated  HN03.     Likewise  test  the  solubility  of  the 
metal  in  aqua  regia.     If  reaction  does  not  take  place  in 
the  cold  with  any  of  the  above-mentioned  acids,  apply 
heat.     Notice  the  gaseous  products  formed  in  each  case. 
(Empty  all  mercury  residues  into  the  bottle  labeled  "Mer- 
cury Residues") 

374.  Prepare  sodium  amalgam  by  adding  two  or  three 
small,  dry,  freshly  cut  pieces  of  sodium  to  a  little  dry 
mercury  in  a  porcelain  mortar.     (CAUTION.)     Examine 
the  product.     Does  it  look  any  different  from  mercury  ? 

Divide  the  amalgam  into  two  portions  in  test  tubes. 
To  one  portion  add  water  and  test  the  gas  evolved  with 
a  burning  splinter.  To  the  other  portion  add  strong 
NH4C1  solution.  Have  you  ever  performed  this  latter 
test  before?  What  becomes  of  the  mercury  in  these 
tests  ? 

Mercurous  Compounds. 

375.  Prepare  a  solution  of  mercurous  nitrate  by  treat- 
ing about  half  a  cubic  centimeter  of  mercury  with  a  little 
moderately  strong  HN03  (i  :  i)  in  a  test  tube.     Allow 
the  tube  to  stand  for  some  minutes.     There  should  be 
some  mercury  left  in  the  bottom  of  the  tube.     Dilute 
the  solution  with  10  cc.  of  water  containing  a  drop  of 
concentrated  HNO3  and  use  in  the  tests  to  follow. 


142  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

376.  Try  the  action  of  NaOH  and  NH4OH  on  separate 
portions  of  HgNO3  solution.     Through  another  portion 
pass  H2S  until  precipitation  is  complete. 

377.  To   separate  portions  of  HgNO3   solution  add 
dilute  HC1  and   a   solution   of  NaCl.     Do   these   two 
reagents  precipitate  the  same  compound  ? 

To  one  of  the  tubes  add  an  excess  of  NH4OH.  Save 
the  other  tube  for  another  experiment. 

Mercuric  Compounds. 

378.  Prepare  a  solution  of  mercuric  chloride  (HgCl2) 
by  dissolving  a  globule  of  mercury  in  aqua  regia.     Evapo- 
rate to  small  volume  and  then  dilute  with  15  cc.  of  water. 
Use  the  solution  in  the  tests  to  follow. 

379.  In  separate  test  tubes  try  the  action  of  NaOH 
solution  and   NH4OH  on  portions  of   the  solution  of 
HgCl2  from  the  preceding  experiment.     (Compare  with 
Exp.  376.)     Through  another  portion  pass  H2S  for  a 
time  and  watch  the  various  changes.     Test  the  solu- 
bility of  this  precipitate  in  strong  HC1  and  HN03. 

380.  To  a  few  cubic  centimeters  of  HgQ2  solution  add 
a  solution  of  KI,  a  little  at  a  time,  until  precipitation  is 
complete.     Avoid  an  excess.     Divide  into  two  portions. 

To  one  portion  add  more  KI  solution  until  the  precipi- 
tate just  redissolves.  Now  add  an  equal  volume  of  strong 
KOH  solution.  What  is  the  name  of  the  mixture  thus 
formed  ?  What  use  is  made  of  this  mixture?  (See  Exp. 
33,  page  27.)  Try  its  action  on  a  solution  of  NH4C1. 

381.  Decant  the  supernatant  liquid  from  the  other 
half  of  the  precipitate  of  HgI2  formed  in  Exp.  380,  and 
then  dissolve  in  a  few  drops  of  concentrated  HC1  by  the 
aid  of  heat.     Allow  to  cool.    What  crystallizes  out  ? 


MERCURY  143 

382.  Treat  the  white  precipitate  of  HgCl  saved  from 
Exp.  377  with  a  little  aqua  regia  and  boil  for  a  moment. 
Does  the  precipitate  dissolve?    Why?    Has  there  been 
a  change  in  the  composition  of  the  mercury  compound? 
Dilute  the  solution  and  apply  tests  to  ascertain  whether 
the  solution  now  contains  a  mercurous  or  a  mercuric 
compound. 

383.  Immerse  a  piece  of  clean  bright  copper  foil  in  a 
little  dilute  HgCl2  solution  and  allow  to  stand  quietly 
for  a  time.     Explain  the  phenomenon  observed. 

384.  To  a  little  HgCl2  solution  add  a  solution  of  stan- 
nous  chloride  (SnCfc).     What  happens?    Add  NH4OH 
to  the  mixture.     Does  this  prove  the  presence  of  a  mer- 
curous or  a  mercuric  compound?     How  has  the  SnCl2 
affected  the  mercuric  salt? 

Summary.  Mention  three  tests  by  which  zinc  and 
cadmium  can  be  distinguished.  In  what  respects  does 
mercury  differ  from  both  zinc  and  cadmium  ?  Compare 
the  solubility  of  the  sulphides  of  mercury,  zinc  and  cad- 
mium. Compare  the  solubility  of  mercurous  and  mer- 
curic compounds. 

Problems,  (a)  From  5  tons  of  zinc  ore  assaying  87%  ZnS,  how 
many  kilos  of  99%  zinc  can  be  theoretically  extracted? 

(b)  What  will  be  the  volume  occupied  by  100  cc.  of  mercury  if 
volatilized  at  1000°  and  740  mm.  pressure? 

(c)  From  150  liters  of  a  solution  containing  10%  of  mercuric 
chloride,  how  many  cubic  centimeters  of  mercury  can  be  obtained  ? 


CHAPTER  XVI. 

COPPER,  SILVER  AND  GOLD. 

COPPER  (Cu;  63). 

385.  In  separate  test  tubes  try  the  action  of  dilute 
HC1,  HNOs  and  H2SO4  on  small  pieces  of  copper.      If 
reaction  does  not  take  place  in  the  cold,  try  the  effect  of 
heat.     Repeat  these  tests  using  concentrated  acids  in- 
stead of  dilute. 

Note  the  products  formed  in  each  case  and  write  all 
equations. 

386.  Into   a  solution   of  CuSO4  introduce  a  bright 
piece  of  sheet  iron  or  an  iron  nail.     Allow  to  stand  for 
a  moment;  then  examine  the  iron.     What  change  has 
taken  place? 

Repeat  the  experiment,  using  a  bright  piece  of  alu- 
minum or  zinc.  Do  these  metals  act  the  same  as 
iron? 

387.  Heat  a  piece  of  bright  copper  for  a  moment  in 
the  upper  (oxidizing)  part  of  the  Bunsen  flame.     Does 
the  copper  change  in  appearance?    Heat  the  piece  of 
copper  a  second  time  and,  while  still  hot,  drop  it  into  a 
test  tube  containing  a  few  drops  of  alcohol.     Explain 
all  changes. 

What  would  happen  if  black  copper  oxide  were  heated 
in  a  stream  of  hydrogen  ? 

388.  In  separate  test  tubes  try  the  action  of  solutions 
of  each  of  the  following  compounds  on  CuSO4  solution: 

K4Fe(CN)6,  Na2C03  and  Na2HPO4. 

144 


COPPER  145 

389.  Treat  a  small  amount  of  CuS04  solution  with 
NH4OH,  a  drop  at  a  time,  until  in  excess.     To  a  liter  of 
water  add  two  or  three  drops  of  CuSO4  solution  and  then 
an  excess  of  NH4OH.     What  do  you  conclude  as  to  the 
delicacy  of  this  test  ? 

To  a  beaker  of  water  add  a  drop  or  two  of  CuS04 
solution  and  then  a  few  cubic  centimeters  of  K4Fe(CN)6 
solution.  (See  Exp.  388.)  Which  is  the  more  delicate 
test  for  copper,  NH4OH  or  K4Fe(CN)6? 

390.  Add  NaOH  to  CuS04  solution  in  a  test  tube. 
Notice  the  color  of  the  precipitate.     Now  heat  the  tube 
and  contents  until  the  liquid  boils.     What  change  has 
taken  place  ? 

391.  Pass  H2S  through  a  solution  of  CuS04.     Filter 
and  wash  the  precipitate.     Test  the  solubility  of  small 
portions  of  the  precipitate  in  HC1  and  in  HNOs. 

392.  Add  KI  solution  to  CuS04  solution.     Filter  and 
wash  the  precipitate.     Test  a  portion  of  the  filtrate  by 
adding  a  few  drops  of  CS2.    What  does  this  test  prove? 
What  is  present  in  solution  in  the  nitrate  ? 

Examine  the  precipitate  on  the  filter.  Is  it  soluble 
in  water?  Heat  a  little  on  a  crucible  cover.  Explain 
the  change  which  takes  place. 

393.  Powder  some  crystals  of  copper  sulphate  in  a 
porcelain  mortar.     What  is  the  color  of  the  powder? 
Put  the  powder  in  a  porcelain  evaporating  dish  and 
heat  gently  over  the  Bunsen  flame.     What  change  takes 
place?    Why  must  care  be  used  to  avoid  heating  too 
strongly  ? 

Allow  the  powder  to  cool;  then  treat  with  a  few  drops 
of  water.  Explain  all  color  changes. 

394.  To  a  few  cubic  centimeters  of  CuS04  solution 


146  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

add  the  same  volume  of  a  solution  of  sodium  and  potas- 
sium tartrate  (NaKC4H4O6).  Then  add  NaOH.  Is 
there  a  precipitate  formed?  (Compare  with  Exp.  390.) 
To  the  mixture  add  a  few  cubic  centimeters  of  grape 
sugar  solution  and  heat  to  boiling.  What  happens  ? 

What  is  the  name  of  the  mixture  of  CuSO4,  NaOH 
and  NaKC4H4O6  solutions?  For  what  is  it  a  test? 
Try  the  action  of  the  mixture  on  a  solution  of  cane 
sugar. 

395.  Treat  a  few  cubic  centimeters  of  CuS04  solution 
with  KCN  solution,  a  drop  at  a  time,  until  the  color  dis- 
appears.    Now  pass  H2S  through  a  portion  of  the  solu- 
tion.    (Compare  with  Exp.  391.) 

396.  Precipitate  copper  ferrocyanide  (Cu2Fe(CN)6)  by 
adding  CuSO4  to  K4Fe(CN)6  solution.     (See  Exp.  388.) 
Treat  the  precipitate  with  strong  NaOH  or  KOH.     Ex- 
plain the  action  of  the  latter  on  the  precipitate. 

397.  Treat  about  5  gms.  of  copper  turnings  with  a 
little  aqua  regia  in  a  small  flask  and  boil  vigorously  for 
a  moment.     Add  5  cc.  concentrated  HC1  and  boil  for  5 
minutes.     Allow  to  settle  for  a  moment  and  then  pour 
the  clear  supernatant  liquid  into  a  large  beaker  full  of 
distilled  water.     Notice  the  color  of  the  precipitate. 

Filter  rapidly  and  wash  the  precipitate  with  a  little 
water.  Test  the  solubility  of  small  portions  of  the  pre- 
cipitate in  HC1,  in  water  and  in  NH4OH.  Heat  a  small 
portion  on  a  crucible  cover. 

What  is  the  composition  of  the  white  precipitate 
formed  above?  Have  you  prepared  any  other  cuprous 
compounds  in  the  preceding  experiments  on  copper  ? 


COPPER  147 

DETERMINATION  OF  THE  ATOMIC  WEIGHT  OF  COPPER 
BY  MEANS  OF  THE  SPECIFIC  HEAT. 

(Law  of  Dulong  and  Petit.) 

398.  Apparatus.  Calorimeter  (consisting  of  two 
beakers,  one  inside  the  other,  the  intervening  space 
being  filled  with  cotton  or  wool),  thermometer,  beaker 
of  boiling  water,  balance,  burner  and  ring  stand,  piece 
of  thread,  metal  to  be  determined. 

Determination  of  Specific  Heat.  Carefully  weigh  the 
empty  calorimeter.  Fill  about  two-thirds  full  of  distilled 
water  and  again  weigh.  The  difference  equals  the  weight 
of  the  water,  W.  Carefully  weigh  the  given  piece  of 
copper  and  let  this  weight  be  represented  by  M.  By 
means  of  a  piece  of  thread,  suspend  the  metal  in  boil- 
ing water  for  several  minutes.  The  metal  is  thereby 
heated  to  100°.  Carefully  note  the  temperature  (T) 
of  the  water  in  the  calorimeter.  Take  the  metal 
from  the  boiling  water,  shake  to  free  it  from  adhering 
water,  and  quickly  introduce  into  the  calorimeter.  Stir 
the  water  in  the  latter  constantly,  at  the  same  time 
watching  the  thermometer  to  note  the  maximum  rise  in 
temperature.  Let  the  maximum  temperature  be  repre- 
sented by  T. 

Data: 

Weight  of  water W 

Weight  of  metal M 

Temperature  of  metal 100° 

Initial  temperature  of  water  in  calorimeter.  .  .        T 
Maximum  temperature  of  water  in  calorimeter      T' 

Inasmuch  as  the  heat  lost  by  the  metal  is  just  equal  to 


148  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

the  heat  absorbed  by  the  water,  the  specific  heat  is  found 
by  the  following  formula: 

Specific  Heat  - 


Determination  of  Atomic  Weight.  According  to  the 
law  of  Dulong  and  Petit,  the  specific  heat  multiplied 
by  the  atomic  weight  is  equal  to  the  constant  6.4. 
Therefore  : 

Atomic  Weight  =  -  -    4          - 
Specific  Heat 

Compare  the  atomic  weight  of  copper,  as  determined 
above,  with  that  given  in  the  table  of  atomic  weights  in 
the  Appendix.  Note  error  and  percentage  error. 

Obtain  from  the  instructor  an  unknown  metal  and 
determine  its  atomic  weight  by  the  method  described 
above.  Compare  the  atomic  weight  thus  determined 
with  the  atomic  weight  table  and  draw  your  conclusions 
as  to  the  metal  employed. 


SILVER  (Ag;  108). 

399.  Dissolve  a  silver  coin  in  dilute  HNO3  and  evapo- 
rate the  solution  to  dryness,  but  do  not  heat  the  dry 
powder.  Dissolve  in  distilled  water  and  filter  if  not 
clear.  Why  is  the  solution  green  in  color  ?  Test  a  drop 
of  the  solution  with  a  few  drops  of  NH4OH.  What  does 
this  prove  ? 

Divide  the  solution  into  two  portions.  Into  one  por- 
tion drop  a  few  minute  pieces  of  copper  turnings  and 
allow  to  stand  quietly  for  some  time.  Then  carefully 


SILVER  149 

examine  the  deposit.  Filter  and  wash  the  crystalline 
precipitate  of  silver.  Pick  out  any  particles  of  copper 
which  remain.  Dissolve  the  silver  in  dilute  HNO3, 
avoiding  an  excess,  and  use  this  solution  of  silver  nitrate 
(AgNOs)  in  the  experiments  to  follow. 

400.  To  the  other  portion  of  the  solution  from  the 
above  experiment  add  dilute  HC1  until  precipitation  is 
complete.     Note  the  color  of  the  precipitate.     Heat  to 
boiling,  filter  and  wash  with  a  little  hot  water  contain- 
ing a  drop  of  HNO3. 

Examine  the  precipitate.  Test  the  solubility  of  small 
portions  of  it  in  HC1,  HNO3  and  in  NH4OH.  To  the 
solution  in  the  latter  reagent  add  concentrated  HN03. 

401.  Add  KCN  solution  to  AgN03  solution,  a  little 
at  a  time,  until  the  precipitate  redissolves.     What  com- 
mercial use  is  made  of  such  a  solution  of  silver? 

402.  Try   the   action   of   solutions   of   the   following 
reagents  on  small  portions  of  AgNO3  solution:   Na2C03, 
NH4OH,  K2Cr04,  NaOH  and  Na2HPO4. 

403.  Pass  H2S  through  a  little  AgNO3  solution  and 
test  the  solubility  of  the  precipitate  in  HNO3. 

Put  a  drop  of  (NH4)2S  solution  on  a  bright  silver  coin. 
What  is  the  chemistry  of  the  action  ?  How  can  the  coin 
be  cleaned  ?  Try  concentrated  HC1  on  it. 

404.  To  about  one-fourth  of  a  test  tube  full  of  AgN03 
solution  add  NH4OH  a  little  at  a  time  until  the  pre- 
cipitate which  is  first  formed  redissolves  and  the  solu- 
tion is  alkaline  to  litmus.     Now  add  about  half  the 
volume  of  a  clear  solution  of  sodium  and  potassium  tar- 
trate  (NaKC4H406),  commonly  called  "Rochelle  Salt," 
and  allow  to  stand  quietly  for  a  time. 

405.  To  separate  portions  of  AgN03  solution  add 


150  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

NaCl,  KBr  and  KI  solutions.  Note  the  color  of  each 
precipitate  at  the  moment  it  is  formed. 

Filter  each  separately  and  spread  the  filter  paper  in 
the  sunlight.  After  a  few  moments  examine  the  pre- 
cipitate. Which  has  been  most  noticeably  changed  by 
the  action  of  the  light?  What  use  is  made  of  these 
silver  salts  ? 

Precipitate  a  little  AgCl  and  test  the  solubility  of  the 
precipitate  in  a  solution  of  sodium  thiosulphate  (Na2S2O3) . 
What  use  does  a  photographer  make  of  Na2S2O3  ? 

406.  Fuse  a  little  of  a  mixture  of  AgCl  (from  Exp. 
400)  and  dry  Na2C03  on  a  piece  of  charcoal,  using  the 
reducing  flame  of  the  blowpipe.     Examine  the  product. 

407.  Test  the  solubility  of  the  metallic  globule  thus 
formed  in  hot  and  cold  concentrated  HC1.     Wash  it 
thoroughly  and  test  with  hot  and  cold  H2S04.     Does 
silver  dissolve  in  either  of  these  acids?    Must  the  acid 
be  heated? 

GOLD   (Au;  197). 

408.  In  separate  test  tubes  try  the  action  of  solutions 
of  the  following  substances  on  small  portions  of  gold 
chloride  (AuCls)  solution  acidified  with  H2SO4i  FeSC>4, 
SnCl2  and  oxalic  acid  (H2C2O4).     Allow  to  settle.     Ex- 
amine the  precipitates. 

Mix  the  precipitates  in  a  flask,  heat  to  boiling  and 
filter.  Wash  the  precipitate  with  a  little  hot  water. 
Note  the  color  of  the  precipitate. 

409.  Heat  a  small  portion  of  the  precipitate  (from 
Exp.  408)  on  a  porcelain  crucible  cover.     What  change 
takes  place  ? 

410.  Try    the    solubility    of    small   portions   of   the 


GOLD  151 

powdered  gold  prepared  in  Exp.  408  in  HC1,  HN03, 
H2S04  and  in  aqua  regia.  What  can  you  say  of  the 
solubility  of  gold  ? 

411.  Combine  all  the  gold  left  from  the  preceding 
experiments,  treat  with  aqua  regia  and  evaporate  care- 
fully   to   dryness.     What   is   the   residue?     Now   heat 
strongly — full  force  of  the  Bunsen  flame.     What  change 
takes  place  ? 

Dissolve  in  aqua  regia  and  again  evaporate  almost  to 
dryness.  Take  up  in  about  15  cc.  of  water  and  use  this 
solution  in  the  tests  to  follow. 

412.  In  separate  test  tubes  try  the  action  of  the  fol- 
lowing reagents  on  a  few  drops  of  AuCl3  solution :  NaOH, 
H2S  and  KI.     Do  any  of  these  reactions  resemble  the 
corresponding  reactions  with  copper? 

(Empty  all  gold  residue  and  solutions  into  the  bottle 
labeled:  u  Gold  Waste.") 

Summary.  Compare  the  valences  of  copper,  silver 
and  gold.  Why  are  these  metals  placed  in  the  same 
group  ?  Why  are  they  grouped  with  the  alkalies  ? 

Compare  the  action  of  NaOH  and  of  NH4OH  on  solu- 
tions of  copper,  silver  and  gold. 

What  is  the  relative  stability  of  copper,  silver  and  gold 
compounds  ? 

Problems,  (a)  From  580  gms.  of  crystallized  copper  sulphate, 
dissolved  in  water  and  treated  with  an  excess  of  KI  solution,  how 
many  grams  of  cuprous  iodide  will  be  produced  ? 

(b)  What  is  the  equivalent  weight  of  copper  ?     Of  silver  ?    Of 
gold?    How  much  copper  will  be  required  to  precipitate  all  the 
silver  from  12  liters  of  a  16%  solution  of  AgN03? 

(c)  A  solution  of  gold  chloride  contains  2%  of  gold.     What 
weight  of  crystallized  ferrous  sulphate  will  be  necessary  to  com- 
pletely precipitate  all  the  gold  from  8  liters  of  this  solution? 


CHAPTER  XVII. 

TIN  AND  LEAD. 
TIN  (Sn;  119). 

413.  In  separate  test  tubes  treat  small  pieces  of  tin 
foil  with  dilute  and  concentrated  HC1,  HNO3  and  H2SO4. 
If  reaction  does  not  proceed  at  the  ordinary  tempera- 
ture, apply  heat.     Try  the  solubility  of  tin  in  aqua  regia. 

Stannous  Compounds. 

414.  Prepare  a  solution  of  stannous  chloride  (SnCl2) 
by  treating  several  small  pieces  of  tin  foil  with  15  cc.  of 
concentrated  HC1  in  a  small  flask.     Warm  gently  to 
hasten  the  reaction.     If  all  the  tin  dissolves,  add  a  few 
more  pieces  —  there  should  be  some  tin  left  to  keep  the 
solution  in  the  stannous  condition 

Dilute  with  15  cc.  of  water.     Label  the  solution  "  Stan- 
nous  Chloride,"  and  use  in  the  tests  to  follow. 

415.  To  a  solution  of  stannous  chloride  (SnCl2)  add 
NaOH  solution,  a  little  at  a  time,  until  precipitation  is 
complete.     What  compound  is  formed?     Now  continue 
to  add  NaOH  solution  until  the  precipitate  redissolves. 
What  compound  of  tin  is  now  in  solution?     Is  it  a 
stannous  compound?     For  what  element  is  the  solution 
thus  prepared  a  good  reagent  or  test? 

416.  Acidify  a  solution  of  KMn04  with  HC1;    then 
add  SnCl2  solution.     Repeat,  using  a  solution  of  K2Cr2O7 
instead  of  KMnO4. 

Try  the  action   of  SnCl2  solution  on  a  solution  of 
HgCl2. 

152 


TIN  153 

What  is  the  chemical  behavior  of  SnCl2  in  these  tests  ? 

417.  Pass  H2S  through  a  few  cubic  centimeters  of 
SnCl2  solution  until  precipitation  is  complete.     Note 
the  color  of  the  precipitate.     Filter  and  wash  the  pre- 
cipitate.    Test  the  solubility  of  small  portions  of  the 
precipitate  in  HC1  and  HNO3.     Also  test  the  solubility 
of  small  portions  in  warm  solutions  of  ammonium  sul- 
phide   ((NH4)2S)    and   in   yellow   ammonium   sulphide 
((NH4)2SJ.     To  the  solution  in  the  latter  reagent  add 
HC1;  note  the  color  of  the  precipitate. 

(Save  a  portion  of  the  precipitate  produced  by  H2S  on 
SnCl2  solution  for  comparison.) 

Stannic  Compounds. 

418.  To  a  few  cubic  centimeters  of  stannic  chloride 
(SnCl4)  solution  add  NaOH,  a  little  at  a  time,  until  in 
excess.     What  compound  is  now  in  solution  ? 

419.  Try  the  action  of  SnCl4  solution  -on  a  solution 
of  HgCl2.     Also  try  it  on  K2Cr2O7  solution.     Is  SnCl4  a 
reducing  agent  ?    Why  ? 

420.  Treat  about  a  gram  of  tin  foil  with  2  or  3  cc. 
of  concentrated  HNOs  in  an  evaporating  dish  under  the 
hood.     When  reaction  ceases,  warm  gently  to  drive  off 
excess  of  HNOs  and  dry  the  powder.     Mix  the  dried 
powder  with  about  twice  its  bulk  of  KCN  on  a  piece 
of  charcoal.     (CAUTION!    Do  not  handle  KCN  with  the 
hands.}     Under  the  hood,  fuse  the  mixture  on  the  char- 
coal, using  the  reducing  flame  of  the  blowpipe.     Heat 
for  several  minutes. 

Cut  away  the  charcoal  where  the  mixture  was  heated, 
grind  it  in  a  mortar,  and  treat  with  water  to  wash  away 
the  particles  of  charcoal.  What  is  left?  Is  it  tin? 


154  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

How  can  it  be  tested  for  tin?  Try  dissolving  it  in  con- 
centrated HC1  and  adding  to  a  solution  of  HgCl2.  If  a 
white  precipitate  is  formed,  what  is  indicated  ? 


LEAD  (Pb;  207). 

421.  In  separate  test  tubes  treat  small  pieces  of  metal- 
lic lead  with  concentrated  and  dilute  HC1,  HN03  and 
H2S04.    Also  try  the  solubility  of  lead  in  HC2H302  and 
in  aqua  regia. 

422.  Heat  a  little  litharge  (lead  oxide,  PbO)  on  char- 
coal with  the  reducing  flame  of  the  blowpipe.     Examine 
the  metallic  globule  of  lead.     Repeat,  using  a  sample  of 
paint  base  to  ascertain  if  it  is  white  lead. 

423.  Immerse  a  strip  of  zinc  in  a  solution  of  lead 
acetate  (Pb(C2H3O2)2)  and  allow  to  stand  quietly.     Can 
you  explain  the  action?    What  is  the  relative  position 
of  these  two  elements  in  the  electrochemical  series  of  the 
elements  ? 

424.  To  separate  portions  of  a  solution  of  Pb(C2H3O2)2 
add  solutions  of  K2Cr04,  Na2C03,  NaCl  or  HC1,  dilute 
H2S04  and  KI.     (See  Exp.  425.) 

425.  Heat  the  tubes  containing  the  precipitates  of 
PbCl2  and  PbI2  from  the  previous  experiment.     What 
happens?     Now  let  the  tubes  stand  quietly  and  cool. 
Note  the  crystals  formed. 

426.  Try  the  action  of  concentrated  HC1  on  lead 
oxide  (PbO)  and  on  lead  dioxide  (PbO2).     Notice  any 
gaseous  products.     Explain  the  difference  in  the  two 
reactions. 

427.  To  a  portion  of  Pb(C2H302)2  solution  add  NaOH, 
a  little  at  a  time,  until  the  precipitate  which  is  at  first 


LEAD  155 

formed  redissolves.     What  two  compounds  have  been 
made  in  this  test  ? 

Divide  the  solution  into  two  equal  portions.  To  one 
portion  add  a  little  HC1  to  cause  reprecipitation. 

428.  To  the  other  part  of  the  solution  formed  in  the 
previous  experiment  add   a   freshly  prepared  solution 
of  NaClO.     What  is  the  composition  of  the  precipi- 
tate? 

429.  Treat  a  little  red  lead  (lead  tetroxide,  Pb304) 
with  dilute  HNO3.     Examine  the  residue.     Dilute  the 
mixture  with  water  and  filter.     Test  the  filtrate  to  as- 
certain if  it  contains  lead. 

430.  Heat  a  little  NaOH  solution  with  a  trace  of 
Pb02.     What  happens?    What  compound  is  formed? 
Is  it  the  same  as  the  compound  formed  in  the  first  part 
of  Exp.  427? 

431.  Pass   H2S   through   a  solution  of   Pb(C2H3O2)2 
until  precipitation  is  complete.     Test  the  solubility  of 
the  resulting  sulphide  in  HN03  and  in  (NH4)2S  solu- 
tion. 

Summary.  To  what  group  of  elements  do  tin  and 
lead  belong?  What  is  the  characteristic  valence  of  the 
elements  of  this  group?  What  other  metallic  element 
belongs  to  this  group?  Why  do  we  not  experiment 
with  this  element  in  the  laboratory? 

Make  a  table  of  all  the  oxides  of  tin  and  lead,  arranging 
them  according  to  the  oxygen  content. 

Problems,  (a)  How  many  grams  of  HgCl  can  be  precipitated 
by  120  cc.  of  a  15%  SnCl2  solution  from  an  excess  of  HgCl2  solu- 
tion? 

(b)  How  many  cubic  centimeters  of  55%  HNO3  will  be  required 
to  oxidize  to  SnCl4  all  the  SnCl2  in  370  cc.  of  a  12%  solution? 


156  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

(c)  How  many  pounds  of  iron  will  be  required  to  reduce  1600 
Ibs.  of  galena  containing  94%  PbS  ? 

(d)  By  oxidizing  8  kilos  of  metallic  lead  to  Pb3O4  and  treating 
this  product  with  HNOa,  how  many  kilos  of  lead  dioxide  can  be 
produced  ? 


CHAPTER  XVIII. 

ALUMINUM  AND  CHROMIUM. 

ALUMINUM   (Al;  27). 

432.  Test  the  solubility  of  metallic  aluminum  in  the 
various  mineral  acids,   both  dilute  and  concentrated. 
Heat  if  necessary. 

Likewise  try  the  action  of  boiling  KOH  or  NaOH  on 
aluminum.  Identify  the  gaseous  products  formed  in 
each  case. 

433.  Test  the  action  of  NH4OH  on  a  solution  of  alu- 
minum.    Add  excess  of  NH4OH  and  then  heat  to  boil- 
ing.   Allow  to  stand  for  a  moment.    Examine  carefully. 

434.  To  a  solution  of  aluminum  sulphate  (A12 (804)3) 
add  NaOH,  a  little  at  a  time,  until  in  decided  excess. 
Test  the  solution  thus  formed  as  follows: 

Heat  a  portion  to  boiling;  allow  to  stand  for  a  moment. 
(Compare  with  Exp.  433.) 

To  a  second  portion  add  strong  NH4C1  solution  and 
heat  gently.  Is  a  gas  evolved?  Test  with  moist  tur- 
meric paper  and  by  the  odor.  What  is  the  precipitate  ? 

435.  Try  the  action  of  Na2COs  solution  on  a  solution 
of  an  aluminum  salt.     Compare  the  precipitate  with 
that  obtained  by  means  of  NH4OH  and  NaOH. 

Filter  and  wash  the  precipitate.  Test  with  HC1  to 
ascertain  if  it  is  a  carbonate. 

Try  the  action  of  a  solution  of  (NH4)2S  on  a  solution 
of  aluminum.  What  is  the  composition  of  the  precipi- 

157 


158  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

tate?  How  can  this  be  explained?  What  action  does 
H2S  have  on  aluminum  solutions?  How  can  aluminum 
sulphide  be  made  ? 

436.  Place  a  little  aluminum  oxide  on  a  piece  of  char- 
coal, moisten  with  a  drop  of  cobalt  nitrate  (Co(NO3)2) 
solution,  and  heat  strongly  before  the  blowpipe.     Notice 
the  color  of  the  product.      This  is  a  good  test  for 
aluminum. 

437.  Treat  a  little  aluminum  solution  with  a  solution 
of  Na2HPO4.     What  is  the  composition  of  the  white 
precipitate  ?     Test  its  solubility  in  acids. 

438.  Preparation  of  Alum  from  Clay.     Under  a  good 
hood  treat  35  gms.  of  clay  with  10  cc.  of  concentrated 
H2SO4  in  an  evaporating  dish.     Heat  the  mixture  gently 
for  about  20  minutes.     Allow  to  cool.     Then  transfer 
the  mixture  to  a  beaker  containing  50  cc.  of  water. 
Rinse  out  the  dish  with  a  little  of  the  water  and  add  the 
washings  to  the  beaker. 

Heat  the  mixture  in  the  beaker  almost  to  boiling  and 
add  about  3  gms.  of  iron  filings.  Keep  warm  for  10 
minutes. 

Dissolve  8  gms.  of  crude  (NH4)2S04  in  the  mixture  and 
filter  while  still  hot.  Evaporate  the  clear  filtrate  to  one- 
half  its  volume  and  set  aside  to  crystallize.  When  cold, 
filter  and  wash  the  crystals  with  a  few  cubic  centimeters 
of  cold  water. 

Taste  the  crystals.  Dissolve  one  in  water  and  test 
the  solution  for  aluminum.  Explain  this  experiment  in 
full. 


ALUMINUM  159 

DETERMINATION  OF  THE  NUMBER  OF  MOLECULES  OF 
WATER  OF  CRYSTALLIZATION  IN  ALUM. 

(Quantitative.) 

439.  In  the  determination  of  the  number  of  molecules 
of  water  of  crystallization  in  alum,   the  temperature 
should  not  be  allowed  to  go  above  200°.      The  crucible 
must  not  be  heated  on  a  triangle  with  the  open  flame 
but  should  be  imbedded  in  a  sand  bath.      The  tem- 
perature is  thereby  more  easily  controlled. 

Aside  from  the  manner  of  applying  the  heat,  the  de- 
termination is  carried  out  exactly  as  the  determination 
of  the  number  of  molecules  of  water  of  crystallization 
in  gypsum  (Exp.  44,  page  31). 

440.  Aluminum  as  a  Mordant.     To  20  cc.  of  alum 
solution  add  a  slight  excess  of  NH4OH,  heat  to  boiling 
and  filter.     Dissolve   the  precipitate   on   the   filter  in 
dilute  HC2H3O2  and  allow  the  solution  to  run  into  a  clean 
beaker. 

Place  a  few  cubic  centimeters  of  this  solution  in  a  test 
tube  and  boil  for  a  moment.  Allow  to  settle.  Ex- 
plain fully. 

Place  a  piece  of  white  cotton  cloth  in  a  beaker  of  dye 
(alizarin)  and  heat  to  boiling.  Then  wash  the  cloth, 
wring  out  the  excess  of  water  and  allow  to  dry. 

Saturate  a  second  piece  of  white  cotton  cloth  in  the 
remainder  of  the  aluminum  acetate  solution  previously 
prepared.  Squeeze  out  the  excess  of  solution,  introduce 
into  the  beaker  of  dye  and  heat  to  boiling  for  a  moment. 
Then  wash  the  cloth  thoroughly,  wring  out  and  dry  as 
before.  Compare  the  two  pieces  of  cloth.  Why  is 
thexe  a  difference  in  color  ? 


160  EXPERIMENTS  IN  GENERAL  CHEMISTRY 

441.  Mordants.     Saturate  one  piece  of  cloth  with  a 
solution  of  alum,  a  second  with  a  solution  of  FeCls  and 
a  third  with  a  solution  of  alum  containing  a  few  cubic 
centimeters   of   FeCls    solution.     The   pieces   of   cloth 
should  be  marked  in  some  way  so  that  they  can  be  dis- 
tinguished. 

In  separate  beakers  now  saturate  the  three  pieces  of 
cloth  in  dilute  NH4OH.  Then  wash  each  piece  thor- 
oughly, wring  out  as  much  of  the  liquid  as  possible, 
and  heat  them  together  for  10  minutes  in  a  beaker  of  the 
dye.  Rinse,  wring  out  and  compare  the  colors. 

(If  the  results  of  this  experiment  are  approved  by  the 
instructor,  the  three  pieces  of  cloth  may  be  trimmed  and 
glued  in  the  notebook.) 

442.  Cement.     Mix   a   few   grams   of   cement    with 
enough  water  to  form  a  thick  paste.     Place  on  a  piece  of 
glass  and  mold  into  a  thin  pat  by  means  of  a  spatula 
or  a  knife  blade.     Allow  to  stand  several  days  to  harden. 

What  is  the  explanation  of  the  hardening  of  cement? 
What  are  the  two  essential  compounds  in  cement  ? 

EXAMINATION  OF  CEMENT.* 

(Qualitative.) 

443.  Silica.     Dissolve  about  a  gram  of  cement  in 
dilute  HC1  in  an  evaporating  dish  and  evaporate  to 
dryness.     Moisten  the  residue  with  a  few  drops  of  dilute 
HC1,  add  20  cc.  of  water  and  warm  gently.     Filter  and 
wash  the  precipitate.     Test  the  precipitate  for  silica  by 
tests  described  in  Exps.  235  and  236.     (Use  the  filtrate 
in  the  next  test.) 

*  The  reagents  used  in  these  tests  must  be  chemically  pure. 


ALUMINUM  l6l 

Alumina  and  Iron.  To  the  filtrate  from  the  preceding 
test  add  NH4OH  until  alkaline,  heat  to  boiling  and 
filter.  What  is  the  white  gelatinous  precipitate  ?  Dis- 
solve a  bit  in  NaOH  and  add  NH4C1  solution.  What 
does  this  prove  ? 

Dissolve  a  portion  in  dilute  HC1,  add  a  drop  of  con- 
centrated HNO3  and  heat  to  boiling.  Cool  under  the 
faucet;  then  add  a  drop  or  two  of  KCNS  solution. 
What  does  this  prove  ?  (Use  the  filtrate  in  the  next  test.) 

Lime.  To  the  filtrate  from  the  preceding  test  add  a 
few  drops  of  NH4OH.  Heat  to  boiling  and  add  an  excess 
of  a  solution  of  (NH4)2C2O4.  Boil  for  a  moment;  filter 
hot  and  wash  with  a  little  hot  water  containing  a  drop 
of  NH4OH. 

Place  the  filter  paper  containing  the  precipitate  in  a 
clean  porcelain  crucible  and  heat  gently  until  dry,  and 
then  strongly  until  the  paper  is  burned.  Now  cover  the 
crucible  and  ignite  in  the  blast  flame.  Allow  to  cool. 
Then  remove  the  cover  and  treat  the»  residue  with  a  few 
cubic  centimeters  of  water.  Filter  through  a  small 
filter  and  test  the  solution  with  turmeric  paper  and  with 
Na^COs  solution. 

(Save  the  filtrate  from  the  precipitation  with  (NH4)2- 
C2O4  for  use  in  the  next  test.) 

Magnesia.  Evaporate  the  filtrate  from  the  above 
test  to  small  volume  (10  or  15  cc.).  Then  add  an  excess 
of  NH4OH  and  (NH^HPO,  solution.  Allow  to  stand 
several  minutes.  What  is  the  composition  of  the  pre- 
cipitate ?  What  compound  would  be  formed  if  this 
precipitate  were  strongly  ignited? 

(Use  the  filtrate  from  this  test  in  the  test  for 
alkalies.) 


1 62  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

Alkalies.  Evaporate  the  filtrate  from  the  preceding 
test  to  dryness  and  heat  gently  to  drive  out  all  am- 
monium salts  (white  fumes).  Cool;  treat  with  two  or 
three  drops  of  dilute  HC1  and  test  for  Na  and  K  by 
means  of  the  flame  test  or  with  the  spectroscope. 

Sulphuric  Acid.  To  test  for  sulphuric  acid  dissolve  a 
small  amount  of  the  cement  in  dilute  HC1,  filter,  and  to 
the  clear  filtrate  add  a  few  drops  of  BaCl2  solution.  If 
the  white  precipitate  is  insoluble  in  concentrated  HNO3, 
it  is  BaS04  —  showing  the  presence  of  SO3  in  the  cement. 

CHROMIUM  (Cr;  52). 

444.  Chromic    Compounds.     To    a    solution    of    a 
chromic  compound  as  chromic  sulphate  (Cr2(SO4)3),  or 
chromic  chloride  (CrCl3),  add  NaOH  solution,  a  little 
at  a  time,  until  in  excess.     Explain  all  changes. 

Divide  the  solution  into  two  portions  and  heat  one  to 
boiling.  Allow  to  stand  for  a  moment;  then  compare 
with  the  other  portion. 

445.  Repeat  the  previous  experiment,  using  NH4OH 
instead  of  NaOH.     Note  all  changes  and  compare  with 
Exp.  444. 

446.  To  a  solution  of  a  chromic  salt  add  a  strong 
solution  of  Na2C03.     What  gas  is  evolved?    Why  is  it 
evolved?     Compare  with  Exp.  435. 

447.  Mix  5  gms.  of  powdered  potassium  dichromate 
(K2Cr2O7)  with  i  gm.  of  flowers  of  sulphur  and  intro- 
duce the  mixture  into  a  porcelain  crucible.     Heat  in  the 
flame  of  the  blast  lamp  for  10  minutes. 

Cool,  boil  the  residue  with  water,  filter  and  dry  the 
green  powder  left  on  the  filter.  Test  its  solubility  in 
acids. 


CHROMIUM  163 

448.  Oxidation  of  Chromic  Compounds  to  Chromates. 
Add  NaOH  to  a  CrCl3  solution  until  the  precipitate 
which  is  first  formed  redissolves.     Then  add  about  an 
equal  volume  of  bromine  water  and  heat  to  boiling. 
Does    the    solution    change    in    color  ?    Acidify    with 
HC2H3O2.     Test  a  portion  with  a  solution  of  BaCl2. 

449.  Make  an  intimate  mixture  of  i  gm.  of  finely 
powdered  chrome  iron  ore  (FeCr2O4),  4  gms.  dry  Na2CO3 
and  2  gms.  KN03.     Fuse  the  mixture  in  an  iron  crucible 
over  the  blast  lamp.     Allow  to  cool,  extract  the  melt 
with  hot  water  and  filter.     Neutralize  the  solution  with 
HC2H302  and  test  a  portion  for  chromates  by  adding 
AgN03  solution. 

450.  Fuse  a  pinch  of  the  green  residue  from  Exp.  447 
on  platinum  foil  with  a  little  Na2CO3.     Why  is  this  a 
good  test  for  chromium?     Dissolve  in  a  drop  of  water 
and  test  as  in  Exp.  449. 

451.  Chromates.     To  separate  portions  of  a  solution 
of   K2CrO4  add  solutions  of   Pb(C2H3O2)2,  BaCl2   and 
AgNO3.     Divide  each  precipitate  into  two  portions  and 
test  the  solubility  in  HNO3  and  in  HC2H3O2. 

452.  Add  a  few  drops  of  concentrated  H2S04  to  a 
solution   of   K2Cr04.     What   change   in   color   do   you 
notice  ?     Repeat  the  experiment,  using  some  other  acid. 

To  a  solution  of  potassium  dichromate  (K2Cr207)  add 
NaOH  solution.  What  change  takes  place  ?  Are  these 
two  reactions  oxidation  reactions?  Are  they  reducing 
reactions  ? 

453.  Chromic  Acid.     To  a  few  cubic  centimeters  of 
K2Cr2O7  solution  add  about  an  equal  volume  of  con- 
centrated H2SO4.     Allow  to  stand  quietly  until  cool. 
What  is  the  nature  of  the  red  crystals  formed?    Pour 


[64  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

off  the  supernatant  liquid  and  test  the  solubility  of  the 
crystals  in  water. 

454.  Reduction  of  Chromates.    To  separate  portions 
of  a  chromate  solution  acidified  with  HC1  add  solutions 
of  H2S  and  SnCl2.     Describe  each  change  and  write  all 
equations. 

To  another  portion  of  a  chromate  solution  acidified 
with  H2SO4  add  a  few  cubic  centimeters  of  alcohol. 
Heat  to  boiling;  note  odor  and  change  in  color.  Ex- 
plain fully  and  write  an  equation  to  represent  the  re- 
action. 

455.  To  a  test  tube  half  full  of  water  add  a  few  drops 
of  K2CrO4  solution  and  a  little  dilute  H2SO4.     Now  add 
about  a  half -inch  layer  of  ether  and  then  a  few  cubic 
centimeters  of  hydrogen  peroxide.     Agitate  slightly  and 
allow  to  stand  quietly  for  an  instant.     Examine  the 
color  of  the  solution  and  of  the  ether.     After  10  minutes 
again  examine  the  colors.     Note  and  explain  all  changes. 
This  is  a  good  test  for  a  chromate  and  for  H202. 

456.  Chrome  Alum  (K2Cr2(SO4)4.24H2O).     Dissolve 
100  gms.  of  K2Cr207  in  warm  water.     Add  85  cc.  of  con- 
centrated H2SC>4  carefully  and  allow  to  cool  to  about  30°. 
Add  alcohol  slowly  until  there  is  no  further  action,  being 
careful  to  keep  the  temperature  down  while  doing  so. 
When  a  further  addition  of  alcohol  does  not  cause  a 
reaction  to  take  place,  filter,  concentrate  and  set  aside 
to  crystallize. 

Summary.  What  is  the  usual  color  of  chromic  com- 
pounds? What  is  the  color  of  most  chromates?  Of 
dichromates  ?  What  is  the  valence  of  chromium  in  each 
of  these  three  series  of  compounds?  Has  chromium 
any  other  valence  ? 


CHROMIUM  165 

Mention  several  ways  in  which  chromium  (i)  differs 
from  and  (2)  resembles  aluminum. 

Problems,  (a)  How  many  cubic  centimeters  of  H2S  gas  at 
75°  F.  and  788  mm.  pressure  will  be  required  to  completely  reduce 
1080  cc.  of  an  acid  solution  containing  16%  K2Cr2O7  and  having 
a  specific  gravity  of  1.112  ? 

(b)  From  i  ton  of  chrome  iron  ore  containing  92%  actual  chro- 
mite  (FeteCM ,  what  weight  of  chrome  yellow  can  be  made  ? 

(c)  A  certain  grade  of  Arkansas  bauxite  contains  65%  A12O3. 
How  many  pounds  of  crystallized  potassium  alum  is  it  possible  to 
make  from  1500  Ibs.  of  this  ore? 


CHAPTER  XIX. 
MANGANESE  (Mn;  55). 

457.  •  Manganous  Compounds.     Treat  separate  por- 
tions of  a  solution  of  manganous  sulphate  (MnSO4)  or 
manganous  chloride  (MnCl2)  with  solutions  of  NaOH, 
Na2CO3  and  (NH4)2S.     Does  the  precipitate  formed  by 
NaOH  change  upon  standing?     Test  the  solubility  of 
the  (NH4)2S  precipitate  in  HC1. 

458.  Treat  a  little  Mn02  with  concentrated  HC1  and 
heat  gently.     What  gas  is  evolved  ?    Dilute  the  solution 
to  about  twice  its  volume  and  filter.     To  the  clear  fil- 
trate add  NaOH  solution. 

459.  Make  a  borax  bead  containing  a  bit  of  some 
manganese  compound  and  heat  in  the  oxidizing  and 
reducing  flames.     What  colors  are  produced  in   each 
flame? 

460.  Manganates.     Fuse  a  little  Na2CO3  with  a  mere 
speck  of  MnO2  on  a  piece  of  platinum  foil.    Cool.    What 
color  has  been  produced  by  the  manganese  ? 

This  is  a  very  delicate  test  for  manganese. 

To  20  cc.  of  KMn04  solution  in  a  beaker  add 
NaOH  until  strongly  alkaline.  Now  add  alcohol  a 
drop  at  a  time,  stirring  constantly,  until  the  solution  is 
green. 

What  does  the  solution  contain?     (See  Exp.  461.) 

461.  Permanganates.    Treat  a  small  portion  of  the 
green  solution  prepared  in  the  previous  experiment  with 

166 


MANGANESE  167 

a  little  dilute  H2SO4.     Note  change  in  color.     Is  this 
due  to  a  chemical  or  physical  change  ?    Explain. 

Pass  C02  through  a  second  portion  of  the  manganate 
solution  from  Exp.  460.  Explain  the  change. 

462.  Reduction  of  Permanganates.     To  a  few  cubic 
centimeters  of  a  dilute  KMn04  solution  add  a  little 
dilute  HC1  and  then  SnCl2  solution.     Acidify  a  second 
portion  of  KMnO4  solution  with  dilute  H2S04  and  then 
add  SO2  solution. 

In  separate  test  tubes  also  try  the  action  of  H2S, 
and  alcohol  on  portions  of  KMnO4  solution  acidified 
with  H2SO4. 

Why  is  the  KMnO4  solution  decolorized  in  each  case  ? 
What  compounds  of  manganese  are  formed  when  a  per- 
manganate is  reduced  in  the  presence  of  an  acid  ?  Write 
all  equations  involved  in  this  experiment. 

463.  Treat  a  little  KMnO4  solution  with  NaOH  until 
strongly  alkaline.     Now  add  a  few  drops  of  alcohol  and 
agitate.    Note  the  color  of  the  solution.     Add  a  little 
more  alcohol  and  then  warm.     Note  the  change.     What 
is  the  precipitate  ? 

Compare  with  the  previous  experiment.  Is  the  reduc- 
tion of  permanganates  in  alkaline  solution  different  from 
that  in  acid  solution  ?  Show  by  equations  the  number 
of  oxygen  atoms  available  when  KMnO4  oxidizes  (i)  in 
acid  solution  and  (2)  in  alkaline  solution. 

EXAMINATION  OF  WATER  FOR  DISSOLVED  OXYGEN. 
(Qualitative.) 

464.  Completely  fill  a  glass-stoppered  bottle  with  the 
water  to  be  tested.     By  means  of  a  pipette  or  a  piece  of 
glass  tubing,  introduce  about  i  cc.  of  a  solution  of  MnSO4 


1 68  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

into  the  bottle  near  the  bottom.  In  like  manner  add 
i  cc.  of  KOH  solution.  Quickly  stopper  the  bottle  and 
shake  to  thoroughly  mix  the  contents. 

The  KOH  causes  the  precipitation  of  manganous 
hydroxide  (Mn(OH)2)  and  this,  coming  in  contact  with 
the  oxygen  dissolved  in  the  water,  is  oxidized  to  manganic 
hydroxide  (Mn(OH)3).  One  atom  of  oxygen  will  oxidize 
two  molecules  of  Mn(OH)2  to  Mn(OH)3. 

Allow  the  tightly  stoppered  bottle  to  stand  about  5 
minutes,  then  remove  the  stopper  and  quickly  introduce 
about  a  cubic  centimeter  each  of  KI  solution  and  con- 
centrated HC1,  using  a  pipette  or  glass  tube  as  before. 
Quickly  stopper  the  bottle  and  shake  to  mix  the  con- 
tents. 

The  HC1  dissolves  the  Mn(OH)3  and  also  the  excess  of 
Mn(OH)2.  With  the  former,  manganic  chloride  (MnCl3) 
is  formed,  but  this  is  not  stable;  hence  it  immediately 
breaks  down  with  formation  of  MnC^  and  the  liberation 
of  chlorine.  The  chlorine  thus  liberated  reacts  with  the 
KI  solution,  liberating  iodine  (I)  which  produces  a  color 
varying  from  yellow  to  brown,  depending  upon  the 
amount  of  oxygen  which  was  dissolved  in  the  water. 

One  molecule  of  Mn(OH)3  with  HC1  liberates  one 
atom  of  chlorine  and  this  in  turn  liberates  one  atom  of 
iodine.  Therefore,  for  every  atom  of  oxygen  which  was 
originally  contained  in  solution  in  the  water,  two  atoms 
of  iodine  are  liberated.  Write  equations  to  represent 
all  chemical  changes  which  take  place  in  this  experi- 
ment. 

This  experiment  should  be  tried  with  several  different 
waters  and  a  statement  made  as  to  which  contains  the 
most  dissolved  oxygen.  If  the  various  tests  show  about 


MANGANESE  169 

the  same  intensity  of  color,  add  a  few  drops  of  starch 
paste  to  each  bottle  before  drawing  conclusions. 

STANDARDIZATION  OF  A  POTASSIUM  PERMANGANATE 

SOLUTION  AND  DETERMINATION  or  IRON  BY 

TITRATION. 

(Quantitative.) 

465.  Carefully  clean  a  piece  of  iron  wire  by  means  of 
sand  paper.  Accurately  weigh  out  about  0.2  gm.  of  the 
clean  wire  and  introduce  it  into  a  clean  flask.  Add  a 
little  sodium  carbonate  to  the  flask  and  then  about  40  cc. 
of  dilute  H2SO4.  Place  a  funnel  in  the  flask  to  serve  as  a 
sort  of  stopper  (Fig.  36)  and  then  heat  gently  until  all 
iron  is  dissolved. 

Obtain  a  supply  (about  50  cc.)  of  "Standard  Per- 
manganate  Solution"    from   the   stock   bottle.     Rinse 
out  a  burette  with  small  amounts  of  this 
solution   and   then   fill   the   burette   and 
clamp  it  in  position.     (See  Fig.  14,  page 

56.) 

Carefully  read  the  level  of  the  solution 
in  the  burette.  Now  remove  the  funnel 
from  the  flask  containing  the  iron  solu- 
tion, and  allow  permanganate  solution  to 
run  into  the  flask  a  drop  at  a  time,  shak- 
ing the  flask  gently  after  each  addition.  FlG 
The  titration  is  complete  (all  iron  is  com- 
pletely oxidized  by  the  permanganate)  when  a  drop  of 
the  permanganate  solution  finally  produces  a  faint  pink 
color.  Read  the  burette  accurately.  How  many  cubic 
centimeters  of  permanganate  solution  did  it  require  to 
oxidize  the  iron  ? 


170  EXPERIMENTS   IN   GENERAL  CHEMISTRY 

Weigh  out  a  second  and  a  third  sample  of  iron  wire, 
dissolve  in  acid  and  titrate  as  before.  Compare  the 
results  from  the  three  titrations  and  calculate  from  each 
result  how  much  iron  each  cubic  centimeter  of  the  per- 
manganate solution  is  equivalent  to.  If  the  results  are 
fairly  close,  take  the  average  of  the  three  determinations. 
This  result,  expressed  as  the  amount  of  iron  equivalent 
to  i  cc.  of  the  permanganate  solution,  is  called  the  iron 
factor  of  the  solution. 

Obtain  from  the  instructor  solutions  of  iron  in  which 
you  are  to  determine  the  exact  amount  of  iron  by  titra- 
tion  with  the  " standardized"  permanganate  solution. 
The  number  of  cubic  centimeters  of  the  permanganate 
solution  required,  multiplied  by  the  iron  factor  of  the 
solution,  gives  the  total  amount  of  iron  present. 

Summary.  How  many  series  of  salts  has  manganese  ? 
Which  of  these  are  common?  Which  series  is  the  least 
common?  Give  formulas  of  the  oxides  of  manganese 
and  underline  once  those  distinctly  basic  in  character 
and  underline  twice  those  distinctly  acid  in  character. 

What  tests  could  you  apply  to  distinguish  between 
manganese  and  chromium  compounds  ? 

Problems,  (a)  How  many  grams  of  KMnO4  can  be  reduced 
by  75  gms.  of  90%  alcohol  (i)  in  acid  solution  and  (2)  in  alkaline 
solution  ? 

(b)  How  many  grams  of  crystallized  ferrous  sulphate  can  be 
oxidized  by  18  gms.  of  KMnO4  in  acid  solution? 

(c)  How  many  grams  of  crystallized  ammonium  manganous 
sulphate  ( ?)  can  be  prepared  from  the  waste  liquor  in  a  chlorine 
generator  in  which  680  liters  of  chlorine  have  been  prepared  by 
the  action  of  MnO2  with  an  excess  of  HC1? 


CHAPTER  XX. 

IRON,  COBALT  AND  NICKEL. 

IRON  (Fe;  56). 

466.  Test  the  solubility  of  iron  in  the  various  mineral 
acids.     Also  try  the  action  of  fuming  nitric  acid  on  iron 
by  dipping  a  piece  of  sheet  iron  into  it. 

Ferrous  Compounds. 

467.  To   separate  portions   of   FeS04   solution   add 
NH4OH  and  NaOH.     Observe  the  color  of  the  pre- 
cipitates ;  allow  to  stand  for  a  time  and  note  any  change. 

468.  Try  the  action  of  H2S  and  of  (NH4)2S  solution 
on  separate  portions  of  FeSO4  solution.     Do  they  react 
the  same?     Why? 

Test  separate  portions  of  FeSO4  solution  with  solu- 
tions of  K4Fe(CN)6)  K3Fe(CN)6  and  KCNS.  These  are 
good  tests;  note  carefully. 

469.  To  a  few  cubic  centimeters  of  a  solution  of  FeSO4 
add  Na2CO3  solution.     Note  the  color  of  the  precipitate. 
Filter  rapidly.     Test  a  portion  of  this  precipitate  with 
HC1.     Does  it  effervesce?     Is  the  precipitate  a  car- 
bonate ? 

Allow  the  remainder  of  the  precipitate  to  stand  exposed 
to  the  air.  Does  it  change?  What  is  the  chemistry 
of  this  change? 

470.  In  separate  test  tubes  try  the  action  of  FeSO4 
solution  on  solutions  of  KMnO4  and    K2Cr207  which 
have  been  acidified  with  H2SO4.     What  is  the  action 
of  FeSO4  in  these  cases? 

171 


172  EXPERIMENTS  IN  GENERAL  CHEMISTRY 

471.  Fuse  a  little  sulphur  with  iron  filings.     Allow  to 
cool.    Break  the  tube  and  treat  the  black  substance 
with  HC1.     Note  any  odor  produced.     What  is  the 
composition  of  the  black  compound? 

Treat  a  little  of  the  mineral  "pyrites"  (FeS2)  with 
HC1.  Is  any  odor  of  H^S  noticeable  ?  Why  is  there  a 
difference  ? 

472.  Ferrous  Sulphate   (FeSO^.      Dilute   50  cc.  of 
concentrated  H2S04  by  pouring  it  into  250  cc.  of  dis- 
tilled water.     In  a  large  evaporating  dish  or  casserole, 
treat  50  gms.  of  iron  filings  or  nails  with  the  dilute  acid. 
When  the  action  becomes  slow,  heat  the  dish  gently. 

Concentrate  the  solution  by  boiling;  filter  and  allow 
the  filtrate  to  cool  and  crystallize.  Pour  off  the  mother 
liquor  and  wash  the  crystals  with  a  very  small  amount 
of  water.  Dry  between  filter  papers. 

Concentrate  the  mother  liquor,  add  5  or  10  cc.  of 
strong  H2SO4  and  two  or  three  small  pieces  of  iron  and 
allow  to  stand  for  a  time  to  reduce  any  ferric  sulphate 
to  ferrous.  Filter  concentrate  and  allow  to  crystallize. 


Ferric  Compounds. 

473.  Prepare  a  solution  of  ferric  chloride  (FeCls)  by 
dissolving  a  few  grams  of  iron  in  aqua  regia.     Evaporate 
the  solution  nearly  to  dryness;   then  dilute  with  50  cc. 
of  distilled  water.     Filter  if  necessary.     Use  the  solution 
for  the  ferric  tests. 

474.  To  separate  portions  of  FeCla  solution  add  NaOH 
and  NH4OH.    Are  the  precipitates  soluble  in  excess? 
Are  the  precipitates  the  same  ?     Compare  with  Exp.  467. 

Try  the  action  of  solutions  of  the  following  substances 


IRON  173 

on  separate  portions  of  FeCl3  solution:  K4Fe(CN)6, 
K3Fe(CN)6  and  KCNS.  Make  careful  note  of  the 
precipitates  and  the  colors.  Compare  with  Exp. 
468. 

475.  Treat  a  solution  of  FeCl3  with  a  solution  of 
Na2C03.     Is   a   gas   evolved?     Can    you   identify   it? 
Compare  the  precipitate  obtained  with  that  produced 
by  Na2CO3  on  a  ferrous  solution  (Exp.  469). 

476.  Pass  H2S  through  a  few  cubic  centimeters  of 
FeCla  solution  for  several  minutes.     Do  you  notice  any 
change?     Filter;    compare  the  clear  filtrate  with  FeCla 
solution.     Are  they  the  same  in  appearance?    Try  the 
action  of  a  solution  of  (NH4)2S  on  a  little  FeCl3  solution. 
Compare  with  Exp.  468. 

477.  To  a  few  cubic  centimeters  of  FeCl3  solution  in 
a  beaker  add  Na2CO3  solution  until  nearly  neutral.     If 
too  much  Na^COs  solution  is  added  and  the  solution  is 
basic  to  litmus,  add  dilute  HC1,  drop  by  drop,  until  the 
solution  is  barely  acid. 

To  the  nearly  neutral  but  slightly  acid  solution  thus 
prepared  add  50  cc.  of  water  and  about  a  gram  of 
Na(C2H302)  crystals.  Heat  to  boiling;  then  allow  to 
settle.  What  is  the  precipitate  ?  Is  it  the  same  as  the 
precipitate  formed  in  Exp.  474  ? 

Filter.  Dissolve  the  precipitate  on  the  filter  by  add- 
ing a  few  drops  of  dilute  HC1  and  let  the  solution  run 
into  a  clean  beaker  or  test  tube.  Warm  gently  and 
notice  the  odor. 

478.  Oxidation  of  Ferrous  Compounds  to  Ferric.    To 
a  few  cubic  centimeters  of  FeSO4  solution  add  a  little 
dilute  H2SO4  and  a  drop  of  concentrated  HNO3.     Heat 
to  boiling,  cool,  and  test  with  KCNS  solution. 


174  EXPERIMENTS   IN   GENERAL   CHEMISTRY 

Repeat,  using  KMnO4  solution  instead  of  HNO3. 
Repeat  a  second  time,  using  bromine  water  instead  of 
HN03. 

479.  Ferric  Alum.     Add  23  cc.  of  concentrated  H2SO4 
to  60  cc.  of  distilled  water  in  a  large  evaporating  dish, 
heat  to   100°   and  add  13  cc.  of  concentrated  HNO3. 
Now  add  120  gms.  of  FeSO4  crystals,  a  little  at  a  time, 
waiting  after  each  addition  until  effervescence  moderates. 
When  all  is  dissolved,  add  concentrated  HNO3,  a  little 
at  a  time,  as  long  as  red  fumes  are  evolved.      Heat  the 
solution  to  boiling  for  a  moment  and  dilute  to  300  Cc. 
Filter  if  necessary. 

Heat  to  boiling,  add  40  gms.  of  (NH4)2SO4  and  30  cc. 
of  dilute  H2SO4.  Set  aside  to  crystallize.  Dry  the 
crystals  between  filter  papers  and  preserve  in  tightly 
stoppered  bottles. 

480.  Reduction  of  Ferric  Compounds  to  Ferrous.     Re- 
duce a  few  cubic  centimeters  of  FeCl3  solution  by  adding 
SnCl2  solution.     Test  the  resulting  solution  with  KCNS. 
Also  test  a  portion  with  K3Fe(CN)6  solution.      (Com- 
pare with  Exp.  468.)     What  other  reducing  agents  could 
be  used  instead  of  SnCl2? 

Drop  a  little  zinc  into  a  test  tube  containing  FeCl3 
solution  and  add  a  few  drops  of  concentrated  HC1. 
Allow  to  stand  for  a  moment.  Then  test  portions  for 
ferrous  and  ferric  salts. 

COBALT  (Co;  59). 

481.  To  a  solution  of  cobaltous  nitrate  (Co(NO3)2) 
add  NaOH  until  precipitation  is  complete.     Note  the 
color  of  the  precipitate.     Now  heat  to  boiling  and  add 


COBALT  175 

bromine  water.  Does  the  precipitate  change  in  color? 
What  two  compounds  of  cobalt  have  been  made  in  this 
experiment  ? 

482.  Treat  separate  portions  of  Co(NO3)2   solution 
with  solutions  of  Na2HPO4  and  Na2C03.     Is  the  pre- 
cipitate obtained  with  the  latter  reagent  a  carbonate? 
Test  it. 

483.  Add  NaOH  to  a  little  Co(NO3)2  solution  until 
precipitation  is  complete.     Filter  and  wash  the  precipi- 
tate.    Dissolve  the  precipitate  on  the  paper  by  adding 
3  or  4  cc.  of  KCN  solution  and  let  the  solution  run  into 
a  test  tube.     Pour  the  solution  through  the  filter  twice. 
What  compound  is  now  in  the  solution  ? 

Now  add  a  little  NaOH  and  then  bromine  water. 
Heat.     Explain  all  changes. 

484.  Add    NH4OH    to    a    little    Co(NO3)2    solution. 
What  is  precipitated  ?     Now  add  more  NH4OH  solution 
to  completely  dissolve  the  precipitate.     Pour  the  solu- 
tion into  a  shallow  dish  and  allow  to  stand  exposed  to 
the  air. 

485.  Acidify  a  few  cubic   centimeters  of   Co(NO3)2 
solution  and  pass  H2S  through  for  a  moment.     Is  a 
precipitate  formed?     Now  add  (NH4)2S  solution.     Does 
the  latter  act  like  H2S  on  a  cobalt  solution  ? 

Filter  and  wash  the  precipitate.     Test  the  solubility 
of  a  small  portion  in  cold  HC1. 

486.  Make  a  borax  bead  and  introduce  a  trace  of  the 
precipitate  formed  in   the  previous   experiment.     Try 
the  action  of  the  oxidizing  and  the  reducing  flame  and 

note  color. 

• 

487.  To  a  solution  of  cobalt  chloride  (CoCl2)  add 
concentrated  HC1.     Can  you  explain  the  change?    Is 


176  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

it  a  physical  or  a  chemical  change?    Now  add  excess 
of  water. 

Carefully  evaporate  a  few  cubic  centimeters  of  CoCl2 
solution  to  dryness.  Cool  and  dissolve  the  residue  in 
a  little  alcohol.  Compare  the  solution  with  an  aqueous 
solution. 

488.  Write  on  a  piece  of  pink  or  white  paper,  using 
an  aqueous  solution  of  CoCl2  instead  of  ink.     Allow  the 
writing  to  dry.     Can  you  see  the  writing  ?    Now  warm 
gently  by  holding  at  some  distance  above  a  Bunsen 
flame.     Explain  the  change.     What  use  can  be  made  of 
this  property  of  CoCl2  ? 

489.  To  a  few  cubic  centimeters  of  Co(N03)2  solution 
add  potassium  nitrite  (KNO2)  solution  and  then  acetic 
acid.     Allow  to  stand  for  a  few  minutes.     Examine  the 
precipitate. 

490.  To  a  test  tube  containing  a  little  water  add  a 
drop  or  two  of  some  cobalt  solution  and  a  little  KCNS 
solution.     Now  add  about  a  half-inch  layer  of  amyl 
alcohol-ether  mixture,  shake  and  examine  the  color  of 
the  amyl  alcohol-ether  layer.     What  do  you  conclude 
as  to  the  delicacy  of  this  test  for  cobalt  ? 

NICKEL  (Ni;  58). 

491.  To  a  solution  of  nickel  nitrate  (Ni(NO3)2)  add 
NaOH.     Observe  the  color  of  the  precipitate.     Heat  to 
boiling  and  add  bromine  water.     Compare  with  Exp. 
481. 

492.  Try  the   action    of   solutions    of   Na2C03   and 
Na!jHPO4  on  separate  portions  of   Ni(NO3)2  solution. 
Does  the  Na2CO3  precipitate  a  carbonate?     Test  it. 

493.  Precipitate  Ni(OH)2  by  means  of  NaOH  solu- 


NICKEL  177 

tion.  Filter  and  wash;  then  dissolve  the  precipitate  in 
KCN  solution  as  in  Exp.  483.  To  the  solution  add 
NaOH  and  bromine  water.  Heat  gently.  Compare 
with  Exp.  483. 

494.  Precipitate  NiS  by  adding  (NH4)2S  solution  to 
any  nickel  solution.     Compare  with  Exp.  485.     Try  the 
solubility  of  a  portion  of  the  precipitate  in  cold  HC1. 

495.  Make  a  borax  bead,  using  a  bit  of  the  NiS  precipi- 
tate from  the  preceding  experiment.     Heat  successively 
in  the  oxidizing  and  reducing  flames  and  note  colors 
produced. 

496.  To  a  few  drops  of  any  nickel  solution  in  a  test 
tube  add  a  little  water  and  then  NH4OH  in  excess.     Now 
add  a  few  drops  of  an  alcoholic  solution  of  dimethyl- 
glyoxime.     Repeat  this  test  with  a  cobalt  solution. 

Summary.  What  is  the  best  test  for  Fe"?  For 
Fe'"?  What  is  the  difference  in  the  action  of  sodium 
carbonate  solution  on  ferrous  and  ferric  solutions  ?  How 
many  oxides  has  iron  ?  What  are  the  chief  ores  of  iron  ? 
What  is  the  principle  involved  in  the  reduction  of  iron 
ores? 

Mention  several  tests  by  which  cobalt  and  nickel 
solutions  can  be  distinguished.  What  is  the  usual  color 
of  nickel  compounds?  Of  cobalt  compounds?  What 
nickel  compound  is  a  decided  exception? 

How  can  nickel  and  cobalt  be  distinguished  in  presence 
of  each  other  ? 

Problems,  (a)  Accurately  weigh  a  United  States  nickel.  How 
much  crystallized  nickel  ammonium  sulphate  (?)  can  be  made 
from  this  coin?  (Is  the  coin  pure  nickel?) 

(b)  What  is  the  percentage  composition  of  crystallized  nickel 
sulphate  ? 


178  EXPERIMENTS  IN   GENERAL  CHEMISTRY 

(c)  How  much  15%  solution  of  KN02  will  be  necessary  to  pre- 
cipitate, as  potassium  cobaltinitrite,  all  the  cobalt  in  380  cc.  of  a 
22%  solution  of  cobalt  chloride?     (Specific  gravity  of  the  cobalt 
solution  is  1.122  and  of  the  KNC>2  solution  is  1.084.) 

(d)  To  a  ladle  containing  40  tons  of  molten  steel,  a  workman 
adds  450  Ibs.  of  25%  ferro- vanadium  and  320  Ibs.  of  68%  ferro- 
manganese.     A  chemical  analysis  of  the  steel  thus  produced  will 
show  what  percentage  of  vanadium  and  what  percentage  of  man- 
ganese? 


CHAPTER  XXI. 
PLATINUM  (Pt;  195) 

497.  Place  what  is  left  of  your  platinum  wire  in  a  test 
tube  and  treat  it  with  hot  concentrated  HC1.     Does  it 
dissolve  ?    Wash  the  wire  and  successively  treat  it  with 
H2S04  and  HNO3.     What  can  you  say  of  the  solubility 
of  platinum  ? 

Now  treat  the  wire  with  aqua  regia  and  warm.  When 
it  is  entirely  dissolved,  evaporate  the  solution  to  small 
volume  and  add  a  cubic  centimeter  or  two  of  strong 
NH4C1  solution.  Allow  to  cool.  Filter  off  the  crystals 
and  wash  them  with  a  few  drops  of  alcohol.  Allow  to 
dry.  (Do  not  throw  the  filtrate  away.) 

What  is  the  composition  of  the  crystals  ?  What  com- 
pound have  we  heretofore  studied  which  is  similar  in 
character  ? 

498.  Transfer  the  dried  crystals  to  a  clean  porcelain 
evaporating  dish  and  heat  carefully  for  a  time;    then 
ignite  strongly.     Examine  the  residue  left  upon  cooling. 
What  is  its  composition?    Dissolve  completely  in  a 
little  aqua  regia,  evaporate  to  small  volume  and  then 
dilute  with  20  cc.  of  water.     Use  this  solution  in  the 
following  experiments. 

499.  To  a  portion  of  the  solution  from  Exp.  498  add 
NaOH   and   grape  sugar  solutions.     Heat   to   boiling. 
Allow  to  stand  for  a  moment.     What  is  the  precipitate  ? 

500.  Pass  H2S  through  the  remainder  of  the  solution. 
Filter  and  wash.    Test  the  solubility  of  the  precipitate 
in  yellow  ammonium  sulphide.    Warm  if  necessary. 

179 


APPENDIX. 


CORRECTION   OF    GAS   VOLUMES   FOR   TEM- 
PERATURE  AND    PRESSURE    CHANGES. 

CORRECTION  FOR  TEMPERATURE  CHANGES. 

According  to  the  law  of  Charles,  the  volume  of  a  gas 
varies  directly  as  the  absolute  temperature,  provided,  of 
course,  that  the  pressure  remains  constant.  In  making 
corrections  for  temperature  changes,  it  is  necessary, 
therefore,  first  to  express  the  temperature  as  absolute 
temperature.  In  practically  .all  scientific  work,  the 
centigrade  scale  of  temperature  is  the  one  employed. 

-2££ Centigrade £ «*>' 

0*  Absolute  273°  375* 

FIG.  37. 

On  the  absolute  scale,  the  degrees  are  of  the  same  value 
as  the  degrees  on  the  centigrade  scale,  but  the  absolute 
zero  is  273°  below  the  centigrade  o°.  The  boiling  point 
on  the  centigrade  scale,  100°,  when  expressed  on  the 
absolute  scale  will  therefore  be  100°  +  273°  =  373°.  The 
relation  between  the  absolute  and  centigrade  scales  can 
readily  be  seen  from  the  accompanying  drawing  (Fig.  37) 
in  which  the  points  on  the  centigrade  scale  are  directly 
above  the  corresponding  points  on  the  absolute  scale. 

To  change  centigrade  into  absolute  it  is  necessary 
simply  to  add,  algebraically,  the  given  centigrade  temper- 

181 


182  APPENDIX 

ature  to  273.  For  example,  30°  C.  expressed  on  the  abso- 
lute scale  will  be  the  sum  of  +  30°  and  +  273°  or  303°. 
Likewise,  —  17°  on  the  centigrade  scale  when  changed 
to  absolute  temperature  will  be  found  by  adding  —  17° 
and  +  273°  which  gives  256°. 

Once  the  temperature  is  expressed  on  the  absolute 
scale,  corrections  for  temperature  changes  are  simple. 
If  we  let 

V  =  given  volume, 

V  =  corrected  volume, 

T  —  given  temperature  (absolute  scale), 

Tr  =  new  temperature  (absolute  scale), 

then,  the  pressure  remaining  constant,  the  corrected 
volume  can  be  found  by  the  proportion 

V  :  V  : :  T  :  T' 
or 


Example.  —  If  a  gas  has  a  volume  of  1200  cc.  at  27°, 
what  volume  will  it  occupy  at  17°? 

This  involves  a  lowering  of  the  temperature  ;  hence  the 
new  volume  will  be  smaller.  The  temperatures,  27°  and 
17°,  expressed  on  the  absolute  scale,  are  respectively 
300°  and  290°.  Substituting  these  values  in  the  propor- 
tion previously  given  we  have 

1200  :V  ::  300  :  290 
or 


300 


APPENDIX  183 

CORRECTION  FOR  PRESSURE  CHANGES. 

According  to  Boyle's  law,  if  the  temperature  remains 
constant,  the  volume  of  a  gas  varies  inversely  as  the  pressure. 
In  other  words,  the  greater  the  pressure  exerted  on  a  gas 
the  smaller  will  be  the  volume.  The  pressure  is  measured 
in  millimeters  of  mercury  and  the  " standard"  condition 
of  pressure  is  the  atmospheric  pressure  at  the  sea  level, 
which  is  equal  to  760  mm. 

•In  making  corrections  for  pressure  changes,  let 

V  =  given  volume, 
V  =  corrected  volume, 
P  =  given  pressure, 
P'  =  new  pressure. 

According  to  Boyle's  law  we  have  the  proportion 

V  :  V  : :  P'  :  P 
and  from  this 

v,    VXP 

P' 

Example.  —  At  750  mm.  pressure  a  gas  occupies  a 
volume  of  800  cc.;  what  volume  will  it  occupy  at  600  mm. 
pressure,  the  temperature  remaining  constant? 

V  =8oocc. 

P  —  750  mm. 

P'  =  600  mm. 

The  corrected  volume,  V,  will  then  be  found  by  the 
proportion 

800  :  V  : :  600  :  750, 

and  from  this  we  have 


184  APPENDIX 

It  frequently  happens  that  gas  volumes  are  measured 
over  water,  in  which  case  the  pressure  on  the  gas  is  some- 
what affected  by  the  pressure  of  water  vapor.  This 
pressure  of  the  water  vapor  is  called  the  "  aqueous 
tension,"  and  varies  with  the  temperature  as  shown  in 
Table  IV  on  page  195.  In  solving  problems  involving 
the  volume  of  a  gas  collected  over  water,  it  is  there- 
fore necessary  to  take  as  the  given  pressure,  not  the 
observed  barometric  pressure,  Pt  but  the  barometric 
pressure  minus  the  aqueous  tension  at  the  observed  tem- 
perature, or  P  —  a. 

CORRECTIONS  FOR  TEMPERATURE  AND  PRESSURE 

COMBINED. 

In  correcting  for  temperature  changes  we  evolved  the 
expression 


and  in  correcting  for  pressure  we  likewise  evolved  a 
formula 

V  -  V  x  P 
~P^ 

In  applying  both  corrections  at  once,  we  obtain  the 
expression 

,  _  V  X  T'  X  P 

TXP' 

It  is  advisable  to  solve  problems  by  such  a  formula,  as 
many  times  it  will  be  found  possible  to  cancel,  thus 
greatly  lessening  the  amount  of  multiplication  and 
division. 

Example.  —  To  find  the  volume  which  1800  cc.  of  a 
gas  at  720  mm.  and  47°  will  occupy  if  the  pressure  is 


APPENDIX  185 

raised  to  86 1  mm.  and  the  temperature  is  lowered  to  if. 
Changing  the  temperature  to  absolute,  and  substituting 
in  the  formula  developed  above,  we  have 

y  __  1800  X  287  X  720 
320  X  861 

which,  by  cancellation  and  multiplication,  gives 
V  =  1350  cc. 


CORRECTION  TO  STANDARD  CONDITIONS. 

Standard  conditions  of  temperature  and  pressure  are 
respectively  o°  C.  and  760  mm.;  hence,  in  correcting  gas 
volumes  to  standard  conditions,  T'  =  o°  C.  or  273°  abso- 
lute, and  P'  =  760  mm.  The  volume  at  standard  con- 
ditions, o°  C.  and  760  mm.,  is  usually  expressed  by  F0. 

Example.  —  What  volume  will  210  liters  of  a  gas  at  12° 
and  ^20  mm.  pressure  occupy  when  reduced  to  standard 
condition  ? 

Using  the  formula  heretofore  developed, 

v,  _  V  X  Tf  X  P 

TXP' 
or 

V        V  X  273  X  P 

rX76o 

we  have 

y   _  210  X  273  X  720, 
285  X  760 

and,  by  cancellation,  we  have 

T7        14  X  273  X  18  .  ... 

Fo  =  -  -  -  190.5  +  liters. 

19  X  19 


1 86  APPENDIX 

CHEMICAL   ARITHMETIC. 

The  importance  of  chemical  arithmetic  cannot  be 
overestimated.  In  general,  chemical  problems  involve 
only  simple  mathematical  principles,  though  it  frequently 
happens  that  several  different  principles  are  touched 
upon  in  the  same  problem.  To  illustrate  the  solving 
of  chemical  mathematical  problems,  the  following  type 
problems  have  been  selected.  A  careful  study  of  these 
and  the  underlying  principles  will  enable  the  student  to 
negotiate  successfully  any  ordinary  problem  in  chemical 
mathematics. 

I.  What  amount  of  CaCh  will  be  produced  by  dissolving 
280  gms.  of  CaO  in  HCl  ? 

The  equation  for  the  reaction  is 

CaO  +  2  HCl  =  CaCl2  +  H2O. 

From  this  we  see  that  from  every  molecule  of  CaO,  one 
molecule  of  CaCl2  is  formed.  Furthermore  the  problem 
deals  with  CaO  and  CaCl2  only;  hence  the  other  quanti- 
ties, H20  and  HCl,  may  be  disregarded. 

Substituting  the  sum  of  the  atomic  weights  under  each 
of  the  substances  involved  we  have 

CaO  +  2  HCl  =  CaCja  +  H2O. 
56  no 

This  shows  that  for  every  56  parts  by  weight  of  CaO, 
there  will  be  produced  no  parts  by  weight  of  CaCl2. 
If  56  parts  by  weight  of  CaO  produce  no  parts  by 
weight  of  CaCl2,  the  amount  of  CaCl2  in  grams  that  can 
be  produced  from  280  gms.  of  CaO  can  be  found  by  the 
proportion 

56  :  no  : :  280  :  #, 


APPENDIX  187 

or 

1 10  X  280 
a-      -^-      =55ogms. 

II.  From  2000  pounds  of  marble,  how  many  pounds  of 
lime  can  be  made  ? 

Marble  is  CaCOs;  lime  is  CaO.     The  equation  for  the 
reaction,  together  with  the  atomic  weights  of  the  sub- 
stances involved  in  the  problem,  is  as  follows: 
CaC03  =  CaO  +  C02. 
100  56 

From  100  parts  by  weight  of  CaCOs,  56  parts  by  weight 
of  CaO  can  be  made.  The  amount  that  can  be  made 
from  2000  pounds  of  lime  can  then  be  found  by  the 
proportion 

100  :  56  : :  2000  :  x 
and 

56  X  2000  ,      . .. 

x  =  - —       —  =  1 1 20  pounds  of  lime. 
100 

III.  (a)  From  800  gms.  of  CaCOs,  treated  with  an  excess 
of  acid,  how  many  grams  of  CO2  can  be  made  ? 

Writing  the  equation,  with  the  sums  of  the  atomic 
weights  of  the  substances  involved,  we  have 

CaC03  +  2  HC1  =  CaQ2  +  jCOH-  H2O, 
loo  44 

and  from  the  proportion 

100  :  44   :  :  800  :  x> 
we  have 

44X800 

x  = > 

IOO 

the  weight  of  CO2  liberated. 


1 88  APPENDIX 

(b)  What  volume  would  the  C02  occupy  ?     In  order  to 
find  the  volume  of  CO2  it  is  necessary  to  find  the  weight 
first  and  then  to  divide  this  by  the  weight  of  unit  volume 
of  the  gas.     This  can  be  done  as  follows:  CC>2  has  a 
vapor  density  of  22  (one-half  the  molecular  weight)  or,  in 
other  words,  it  is  22  times  as  heavy  as  an  equal  volume 
of  hydrogen.     But  a  liter  of  hydrogen  weighs  0.0899  g™--, 
hence  the  weight  of  a  liter  of  CO2  equals  22  X  0.0899  g™-- 
By  dividing  the  weight  of  C02  found  under  part  a, 

44  X  800 
100 

by  the  weight  of  a  liter  of  CO2,  we  have  the  expression 
for  V,  the  volume  of  the  gas  in  liters, 

y  _          44  X  800 

100  X  22  X  0.0899 

(c)  What  volume  would  the  CO2   occupy  at  18°  and 
755  mm.  pressure  ?     The  expression  developed  in  b  rep- 
resents the  volume  of  C02  under  standard  conditions, 
o°  C.  and  760  mm.  pressure,  inasmuch  as  the  weight  of 
a  liter  of  hydrogen  (0.0899  Sm-)  under  these  conditions 
was  used. 

To  find  the  volume  which  the  gas  would  occupy  at 
1 8°  and  755  mm.  it  is  necessary  to  find  first  the  volume 
at  o°  and  760  mm.  as  described  above,  and  then  to 
correct  this  for  the  temperature  and  pressure  change. 
The  equation  for  this  correction  is 

v,  _  v  x  r  x  P 

TXP' 

in  which 

,r  44  X  800 


100  X  22  X  0.0899 


APPENDIX  159 

T=      o°(=273°), 

r=  i80(=29i°), 

P  =  760  mm., 
Pf  =  755  mm. 
By  substituting  in  the  formula  above,  we  have 

yf 44  X  800  X  291  X  760 

100  X  22  X  0.0899  X  273  X  755' 

which,  when  solved,  gives  F',  the  volume  of  COz  at  18° 
and  755  mm.  which  can  be  obtained  from  800  gms.  of 
CaC03. 

IV.  In  order  to  prepare  1250  liters  of  hydrogen  at  12° 
and  785  mm.  pressure,  what  weight  of  zinc  will  be  necessary? 
We  cannot  directly  substitute  the  value,  1250  liters, 
in  a  proportion  to  find  weight  of  zinc,  for  this  would  be 
unduly  mixing  weight  and  volume.  The  weight  of  the 
1250  liters  of  hydrogen  at  12°  and  785  mm.  must  be  found 
first.  We  know  the  weight  of  a  liter  of  hydrogen  at 
o°  and  760  mm.  to  be  0.0899  S111-;  but  we  do  not  know  the 
weight  of  a  liter  of  hydrogen  at  12°  and  785  mm.  The 
volume  must,  therefore,  first  be  corrected  to  standard 
conditions,  o°  and  760  mm.,  before  the  weight  can  be 
found.  The  formula  for  this  correction  is 
T/  VXT'XP 

TXP' 

and,  substituting  the  values,  we  have 
1250  X  273  X  785 

285  X  760 

which  expression  gives  the  volume  at  o°  and  760  mm. 
Multiplying  this  by  0.0899,  we  find 

Wt  of  H  =  1250X273  X7&5  X  0.0899 
285  X  760 


1 90  APPENDIX 

By  the  equation 

_Zn  +  2  HC1  =  ZnCl2  +_2H_ 

65  2 

we  find  that  65  parts  by  weight  of  zinc  give  two  parts  by 
weight  of  hydrogen,  and  therefore  the  amount  of  zinc 
necessary  to  produce  the  weight  of  hydrogen  found  above 
can  be  ascertained  by  the  proportion 

65  :  2  : :  x  :  Wt.  of  H, 

in  which  x  is  the  weight  of  zinc  necessary; 
_  65  X  Wt.  of  H 

2 

and  substituting  the  value  found  for  the  weight  of  hydro- 
gen, we  have 

=  65  X  1250  X  273  X  785  X  0.0899^ 
2  X  285  X  760 

which  equals  the  weight  of  zinc  necessary  to  produce 
1250  liters  of  hydrogen  at  12°  and  785  mm.  pressure. 

V.  What  volume  of  30%  H2SO4  (Sp.  Gr.  1.222)  will  be 
required  to  dissolve  884  gms.  of  FezOs? 

By  the  equation 

FeaQg  +  3  H2SO4  =  Fe2(SO4)3  +  3  H2O, 
160  294 

we  find  that  160  gms.  of  Fe203  require  296  gms.  of  actual 
H2S04.  Then  the  amount  required  to  dissolve  884  gms. 
of  Fe203  can  be  found  by  the  proportion 

160  :  294  :  :  884  :  x, 


APPENDIX  IQI 

and  from  this  we  have 

294  X  884 
x  — 

1 60 

This  gives  the  weight  of  actual  H2S04,  but  actual  H2SO4 
means  100%  H2SO4.  It  will,  of  course,  take  consider- 
ably more  dilute  acid  than  100%  acid.  The  amount  of 

actual  acid,    9      -gms.,  is  therefore  only  30%  of  the 

I  DO 

weight  of  30%  H2SO4  needed.     This  weight  may  be 

expressed  by  multiplying  the  expression  by  — ,  thus 

3° 
producing  the  expression 

294  X  884  X  IPO 
160  X  30 

which  equals  the  weight  of  30%  H2S04  necessary. 

But  the  question  asks  for  volume  of  30%  H2S04  in- 
stead of  weight.  The  volume  in  cubic  centimeters  is 
found  by  dividing  the  total  weight  in  grams  by  the 
weight  of  i  cc.  of  the  acid  in  grams.  The  question  now 
arises:  How  much  does  i  cc.  of  30%  H2SO4  weigh? 
But  the  specific  gravity  of  acid  of  this  strength  was  given 
as  1.222.  This  means  that  the  acid  is  1.222  times  as 
heavy  as  an  equal  volume  of  water.  But  i  cc.  of  water 
weighs  i  gm.;  therefore,  i  cc.  of  30%  H2SO4  weighs  1.222 
X  i  gm.  or  1.222  gms. 

By  dividing  the  expression  for  grams  of  30%  H2S04, 
heretofore  developed,  by  1.222,  we  get  the  volume  of  the 
acid  in  cubic  centimeters;  thus, 

IT-  i         .  204  X  884  X  100 

Volume  in  cc.  =  — —  — 

160  X  30  X  1.222 


Ip2  APPENDIX 

But  if  the  volume  is  desired  in  liters,  the  expression  must 
be  divided  by  1000,  inasmuch  as  there  are  1000  cc.  in  a 
liter. 

Volume  in  liters  = 294X884X100 

160  X  30  X  1.222  X  1000 

This,  then,  gives  the  volume  in  liters  of  30%  H2S04 
necessary  to  dissolve  884  gms.  of  Fe203. 

VI.  The  converse  of  the  last  problem  is  equally  simple. 
For  example:  What  amount  of  MgO  can  be  dissolved  by 
4200  cc.  of  43%  #2S04  (Sp.  Gr.  =  1.333)  ? 

The  total  weight  of  acid  is  4200  X  1.333  gms.,  and  if 
the  acid  is  but  43%  pure,  the  total  amount  of  actual 
H2SO4  is 

4200X1.333X43 
100 

By  the  equation 

MgO  +  H2S04  =  MgS04  +  H2O, 
40          98 

we  find  that  98  parts  by  weight  of  actual  H2S04  will 
dissolve  just  40  parts  by  weight  of  MgO.  By  using  the 
above  expression  for  weight  of  actual  H2SO4  and  sub- 
stituting in  the  proportion,  we  have: 

(4*°o  X  i.333X 

\  IOO 

or 

_  4200  X  1.333  X  43  X  40 
98  X  loo 

in  which  x  equals  the  weight  of  MgO  which  can  be  dis- 
solved by  4200  cc.  of  43%  H2S04. 


APPENDIX 


193 


TABLE  I.  —  APPROXIMATE  ATOMIC  WEIGHTS. 


Aluminum  

Al 

27 

Neodymium 

Nd 

14.3 

Antimony  

Sb 

1  2O 

Neon 

Ne 

2O 

Argon         

A 

4O 

Nickel 

Ni 

<;8 

Arsenic     

As 

7$ 

Nitrogen   . 

N 

14. 

Barium     

Ba 

137 

Osmium   . 

Os 

IQI 

Bismuth     

Bi 

208 

Oxygen 

o 

16 

Boron       

B 

II 

Palladium  

Pd 

1  06 

Bromine  

Br 

80 

Phosphorus.  .  .  . 

P 

31 

Cadmium 

Cd 

112 

Platinum 

Pt 

10? 

Caesium 

Cs 

172 

Potassium 

K 

3O 

Calcium 

Ca 

4.O 

Praseodymium 

Pr 

I4.O 

Carbon  

c 

12 

Rhodium  

Rh 

IO3 

Cerium  

Ce 

I4O 

Rubidium  

Rb 

85 

Chlorine 

Cl 

•2  f 

Ruthenium 

Ru 

IOI 

Chromium 

Cr 

C2 

Samarium 

Srn 

I  ^O 

Cobalt 

Co 

gQ 

Scandium 

Sc 

4.4. 

Columbium 

Cb 

Q? 

Selenium 

Se 

70 

Copper              .    .  . 

Cu 

63 

Silicon 

Si 

28 

Erbium          

Er 

166 

Silver 

Ag 

1  08 

Fluorine         

F 

•[Q 

Sodium 

Na 

23 

Gallium   

Ga 

7O 

Strontium  .... 

Sr 

87 

Germanium  

Ge 

73 

Sulphur  

s 

32 

Glucinum  

Gl 

o 

Tantalum  

Ta 

183 

Gold  

Au 

IQ7 

Tellurium  

Te 

127 

Helium  

He 

4 

Terbium  

Tb 

1  60 

Hydrogen  

H 

I 

Thallium  

Tl 

204 

Indium  

In 

114 

Thorium  

Th 

232 

Iodine  

I 

127 

Thulium  

Tu 

171 

Iridium 

Ir 

IQ3 

Titanium 

Ti 

48 

Iron 

Fe 

<?6 

Tin 

Sn 

IIQ 

Krypton 

Kr 

82 

Tungsten 

W 

184 

Lanthanum 

La 

I  30 

Uranium 

u 

238 

Lead 

Pb 

2O7 

Vanadium  

V 

CJ 

Lithium 

Li 

7 

Xenon       

X 

128 

Magnesium    

Mg 

24. 

Ytterbium  

Yb 

173 

Manganese  

Mn 

ri? 

Yttrium  

Y 

89 

Mercury  

Hg 

2OO 

Zinc  

Zn 

65 

Molybdenum 

Mo 

06 

Zirconium 

Zr 

oo 

APPENDIX 


TABLE   II.  — METRIC  WEIGHTS  AND  MEASURES, 

i  meter  (m.)         =      100  centimeters  (cm.) 

=    1000  millimeters  (mm.) 

=  39.37  inches  (in.) 

=  1.094  yards  (yd.) 

i  liter  (1.)  =    1000  cubic  centimeters  (cc.) 

=  33.81  fluid  ounces  (  fl.  oz.) 

=  1.057  quarts  (qt.) 

i  kilogram  (kg.)   =    1000  grams  (gm.) 

=  35.27  ounces  (oz.) 

=  2 . 205  pounds  (Ib.) 

TABLE  III.— SPECIFIC  GRAVITY  AND  MELTING  POINT 
OF  THE  ELEMENTS. 


Element. 

Sp.  Gr. 

M.P. 

Element. 

Sp.  Gr. 

M.P. 

Aluminum 

2  60 

6<4.    < 

M^agnesium 

I    74. 

632  6 

Antimony 

6  62 

42  C—  4.  Co 

Manganese 

7    3Q 

1900 

Arsenic  

5-73 

Mercury  

12  .  CC 

—  30.38 

Barium  

•2  .  7C. 

85 

Molybdenum 

R  6 

OX,  ° 

w.  h. 

Bismuth  

9   8 

268    3 

Nickel 

8  9 

14.  co 

Boron  

2  .  z 

(?) 

Osmium  

22.48 

2  COO 

Bromine  

7  .  It 

—  7.  3 

Palladium  

11.4 

I  COO 

Cadmium 

8  64. 

316 

Phosphorus 

i  81 

4.4.    2 

Caesium 

i  88 

6     s 

26  t; 

Platinum 

21  5 

1770 

Calcium.  . 

i   c8 

760 

Potassium 

o  87 

c8 

Carbon 

Rhodium   . 

12  .  1 

2OOO 

(a)  Diamond  .  . 

3  .  ^2 

Rubidium  

I  •  ^2 

38     <? 

(6)  Graphite... 

2.  3 

Ruthenium  

12.  26 

1800  (?) 

(c)  Amorphous 

i  8 

Selenium 

4.   80 

217 

Cerium 

6  68 

623 

Silicon 

2    30 

(?) 

Chromium. 

6  co 

1^1  < 

Silver   ...      . 

IO    =CO 

0^4 

Cobalt  

8  6 

1800 

Sodium  

o  08 

o<  .6 

Copper 

8  03 

10^4. 

Strontium 

2    ^4 

(?) 

Gallium  

C.Q< 

30.  i1? 

Sulphur 

Germanium  

c.47 

900 

(a)  Rhombic.  . 

2.07 

114.5 

Glucinum  

I  .Q3 

90o(?) 

(6)  Monoclinic 

i  .96 

1  20 

Gold.  .  .    . 

10    32 

103? 

Tellurium  

6.25 

4C2 

Indium  

7    42 

176 

Thallium  

11.85 

290 

Iodine  

4  Q% 

1  1  3-"s 

Titanium  

3.54 

(?) 

Iridium 

22    42 

22OO 

Tin 

7  •  2O 

232  .  7 

Iron 

7  86 

1^87 

Tungsten  

IO.  I 

2900  (?) 

Lanthanum 

6  i 

(?) 

Uranium  

18.7 

w.  h. 

Lead  

II    37 

^22 

Vanadium  

5-5 

(?) 

Lithium 

o  <JQ 

1  80 

Zinc 

7.  1 

41  f.  ? 

APPENDIX 


195 


TABLE  IV.  —  AQUEOUS  TENSION  IN  MILLIMETERS  OF 
MERCURY. 


Temp.  C. 

Pressure. 

Temp.  C. 

Pressure. 

Degrees. 

Degrees. 

0 

4.6 

18 

15-4 

I 

4-9 

iQ 

16.4 

2 

5-3 

20 

17.4 

3 

5-7 

21 

I8.S 

4 

6.1 

22 

19.7 

5 

6-5 

23 

20.9 

6 

7.0 

24 

22.2 

7 

7-5 

25 

23.5 

8 

8.0 

26 

25.0 

9 

8.6 

27 

26.S 

10 

9-2 

28 

28.1 

ii 

9.8 

29 

29.8 

12 

io-5 

30 

31.6 

13 

II  .2 

31 

33-4 

14 

ix  .9 

32 

35-4 

IS 

12.8 

33 

37-4 

16 

13.6 

34 

39-6 

17 

14.4 

35 

41.9 

TABLE   V.  — TABLE   OF  HARDNESS. 
(On  the  Scale  of  10.) 


Agate 

7 

Iron     

A—? 

Aluminum 

2 

Kaolin  

I 

Antimony 

•2      ? 

Lead  

I.  "? 

Apatite        .  .    . 

C 

Lithium  

0.6 

Arsenic 

3? 

Miagnesium 

2 

Barite 

•2      -2 

Magnetite 

6 

Bismuth 

2    ^ 

Meerschaum 

2—7 

Caesium 

O    2 

Opal              

4-6 

Calcite 

Palladium  

4  8 

Cooper 

2    C—  5 

Platinum  

4   3 

Corundum 

Potassium 

O    S 

Diamond. 

IO 

Quartz 

7" 

Feldspar 

6 

Rubidium 

O    3. 

Flint 

7 

Silver     . 

2    <—  2. 

Fluorite 

Sodium  

O   4. 

Glass  

A      C—  6      C 

Steel   

?-8  5 

Gold  

2  t;—  3 

Sulphur  

I  .  <—  2  .  ? 

Graphite  

o  <—  i 

Talc  

I 

Gypsum 

2 

Tin 

i  8 

Halite 

2 

Topaz 

8 

Iridium  .  .  . 

6 

Zinc                      .    .    . 

2    < 

196 


APPENDIX 


TABLE  VI.  — SPECIFIC  GRAVITY  OF  H2SO4  AT  15°  C. 


Per 

Cent 
Acid. 

Sp.  Gr  . 

Per  Cent 
Acid. 

Sp.  Gr. 

Per  Cent 
Acid. 

Sp.  Gr. 

.  I 

.006 

35 

.264 

69 

•603 

2 

.013 

36 

.272 

70 

.615 

3 

.020 

37 

.281 

71 

.627 

4 

.026 

38 

.289 

72 

.638 

5 

•033 

39 

.298 

73 

.650 

6 

.040 

40 

.306 

74 

.662 

7 

.067 

4i 

•315 

75 

.674 

8 

•054 

42 

-324 

76 

.686 

9 

.061 

43 

•333 

77 

.698 

10 

.068 

44 

•  342 

78 

.709 

ii 

•075 

45 

•351 

79 

.721 

12 

.082 

46 

.361 

80 

•732 

13 

.090 

47 

-370 

81 

•743 

14 

.097 

48 

.380 

82 

•754 

IS 

.104 

49 

-389 

83 

.765 

16 

.  112 

50 

•399 

84 

•775 

17 

.119 

5i 

.409 

85 

.784 

18 

.127 

52 

.419 

86 

•793 

!Q 

•135 

53 

.429 

87 

.801 

20 

.142 

54 

•439 

88 

.808 

21 

.150 

55 

•  449 

89 

.814 

22 

.158 

56 

.460 

90 

.820 

23 

.166 

57 

.470 

91 

-825 

24 

•174 

58 

.481 

92 

.829 

25 

.l8l 

59 

.491 

93 

•833 

26 

.190 

60 

.502 

94 

•  836 

27 

.198 

61 

•513 

95 

•839 

28 

.206 

62 

•524 

96 

.841 

29 

.214 

63 

•535 

97 

.844 

3° 

.222 

64 

•546 

98 

•844 

31 

.230 

65 

•558 

99 

.842 

32 

.238 

66 

•569 

IOO 

•839 

33 

.247 

67 

.580 

34 

•255 

68 

•592 

APPENDIX 


197 


TABLE  VII.  —  SPECIFIC  GRAVITY  OF  HC1  AT   15°  C. 


Per  Cent 
Acid. 

Sp.  Gr. 

Per  Cent 
Acid. 

Sp.  Gr. 

I 

.005 

21 

.  106 

2 

.OIO 

22 

.in 

3 

.015 

23 

.116 

4 

.O2O 

24 

.122 

5 

.025 

25 

.127 

6 

.030 

26 

.132 

7 

•035 

27 

•137 

8 

.040 

28 

•143 

9 

•045 

29 

.148 

10 

.050 

30 

•153 

ii 

•055 

31 

.158 

12 

.060 

32 

.163 

13 

.065 

33 

.168 

14 

.070 

34 

•174 

15 

•075 

35 

•179 

16 

.080 

36 

.184 

17 

•085 

37 

.ISQ 

18 

.090 

38 

•195 

19 

•095 

39 

.  2OO 

20 

.IOO 

40 

.205 

APPENDIX 


TABLE  VIII.  —  SPECIFIC  GRAVITY  OF  H3PO4  AT  15°  C. 


Per 
Cent 
Acid. 

Sp.  Gr. 

Per  Cent 
Acid. 

Sp.  Gr. 

Per  Cent 
Acid. 

Sp.  Gr. 

I 

.006 

30 

.189 

59 

I-43I 

2 

.Oil 

31 

.196 

60 

I-44I 

3 

.017 

32 

.203 

61 

1,451 

4 

.023 

33 

.211 

62 

1.461 

5 

.028 

34 

.218 

63 

I.47I 

6 

•034 

35 

.226 

64 

1.481 

7 

.040 

36 

.234 

65 

1.491 

8 

.046 

37 

.241 

66 

I.50I 

9 

.052 

38 

.249 

67 

I.5II 

10 

.058 

39 

•257 

68 

1.522 

ii 

.064 

40 

.  264 

69 

1-532 

12 

.070 

4i 

.272 

70 

1-543 

13 

.076 

42 

.280 

7i 

1-554 

14 

.082 

43 

.288 

72 

1.564 

15 

.088 

44 

.297 

73 

1-575 

16 

.094 

45 

•305 

74 

1.586 

17 

.102 

46 

.313 

75 

1-597 

18 

.107 

47 

.322 

76 

i.  608 

iQ 

.114 

48 

•330 

77 

i  .619 

20 

.I2O 

49 

•339 

78 

1.630 

21 

.127 

50 

.348 

79 

i  .642 

22 

-133 

5i 

•357 

80 

1-653 

23 

.140 

52 

•367 

81 

i  .664 

24 

.147 

53 

-375 

82 

1.676 

25 

•153 

54 

•384 

83 

1.687 

26 

.160 

55 

•393 

84 

1.698 

27 

.167 

56 

•403 

85 

i  .710 

28 

I.I74 

57 

.412 

29 

I.I82 

58 

.421 

APPENDIX 


199 


TABLE  IX. —  SPECIFIC  GRAVITY  OF  HNO3  AT   15  C' 


Per 
Cent 
Acid. 

Sp.  Gr. 

Per  Cent 
Acid. 

Sp.  Gr. 

Per  Cent 
Acid. 

Sp.  Gr. 

I 

.006 

35 

1  .219 

69 

1.418 

2 

.on 

36 

1  .226 

70 

1.423 

3 

.017 

37 

1-233 

71 

1.427 

4 

.023 

38 

1.239 

72 

I-43I 

5 

.028 

39 

.246 

73 

1-435 

6 

.034 

40 

•253 

74 

1-439 

7 

•039 

4i 

•259 

75 

1-443 

8 

-045 

42 

.266 

76 

1-447 

9 

•051 

43 

.272 

77 

1-450 

10 

•  057 

44 

.279 

78 

1-454 

ii 

.063 

45 

.285 

79 

1.458 

12 

.069 

46 

.292 

80 

i  .462 

13 

•075 

47 

.298 

81 

1-465 

14 

.081 

48 

.304 

82 

1.469 

IS 

.088 

49 

•311 

<     83 

1.472 

16 

.094 

50 

.317 

84 

1-475 

17 

.100 

5i 

.323 

85 

1-479 

18 

.107 

52 

•329 

86 

1.481 

iQ 

•113 

53 

•335 

87 

1-485 

20 

.120 

54 

•341 

88 

1.487 

21 

.126 

55 

•347 

89 

1.490 

22 

•133 

56 

•353 

90 

1-493 

23 

.139 

57 

•358 

9i 

1-495 

24 

.148 

58 

.364 

92 

1.498 

25 

•152 

59 

•369 

93 

1.500 

26 

•159 

60 

•375 

94 

1.502 

27 

.165 

61 

.380 

95 

i-5°4 

28 

.172 

62 

-385 

96 

1.506 

2Q 

.179 

63 

•390 

97 

1.508 

3° 

•185 

64 

•395 

98 

1.511 

31 

.192 

65 

.400 

99 

1.516 

32 

.199 

66 

•  405 

IOO 

1-525 

33 

.206 

67 

.409 

34 

.212 

68 

.414 

200 


APPENDIX 


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INDEX. 


A. 

Acids,  bases  and  salts,  53. 

Acidimetry  and  alkalimetry  (quan- 
titative), 55. 

Agencies  of  chemical  change,  i. 

Air,  59- 

Air,  analysis  of  (quantitative),  60. 

Alkaline  earths,  128. 

Aluminum,  157. 

Alum,  water  of  crystallization  in 
(quantitative),  159. 

Ammonia,  62. 

Ammonia,  weight  of  a  liter  of 
(quantitative),  65. 

Ammonium,  125. 

Antimony,  114. 

Aqua  Regia,  71. 

Arsenic,  in. 

B. 

Barium,  134. 

Barium  chloride,  water  of  crystal- 
lization in  (quantitative),  135. 
Bases,  53 . 

Basicity  of  an  acid,  54. 
Bismuth,  115. 
Blowpipe,  75. 
Borax  bead,  74. 
Boron,  102. 
Bromine,  46. 


Cadmium,  140. 
Calcium,  128. 


Carbon,  87. 

Carbon  dioxide,  weight  of  a  liter  of 
(quantitative),  93. 

Carbonic  acid,  91. 

Cement,  examination  of,  160. 

Chemical  arithmetic,  186. 

Chemical  changes,  i. 

Chlorine,  35. 

Chlorine,  oxygen  acids  of,  43. 

Chlorine,  weight  of  a  liter  of 
(quantitative),  38. 

Chromium,  162. 

Coal,  analysis  of  (quantitative), 
88. 

Cobalt,  174. 

Copper,  144. 

Copper,  atomic  weight  of  (quan- 
titative), 147. 

Correction  of  gas  volumes,  181. 

Cyanogen  and  the  cyanides,  95. 

D. 

Definite  proportions,  verification 
of  the  law  of  (quantitative),  20. 

F. 

Fertilizers,  examination  of,  132 
Fluorine,  51. 

G. 

Gold,  150. 

Gypsum,  water  of  crystallization 
in  (quantitative),  31. 


203 


2O4 


INDEX 


H. 

Halogens,  35. 
Hydriodic  acid,  50. 
Hydrobromic  acid,  47. 
Hydrochloric  acid,  41. 
Hydrofluoric  acid,  51. 
Hydrogen,  3. 

Hydrogen  equivalents,  10. 
Hydrogen  peroxide,  33. 
Hydrogen,  reduction  by  (quanti- 
tative), 8. 

Hydrogen  sulphide,  79. 
Hydroxylamine,  64. 


I. 

Iodine,  49. 
Iron,  171. 

Iron,  determination  of,  by  titra- 
tion  (quantitative),  169. 

L. 

Law  of  definite  proportions,  veri- 
fication of,  20. 

Law  of  Dulong  and  Petit,  147. 
Lead,  154. 
Lithium,  118. 

M. 

Magnesium,  137. 
Manganese,  166. 
Marsh  test,  113. 
Mercury,  140. 


N. 


Nickel,  176. 
Nitric  acid,  69. 
Nitrogen,  58. 
Nitrous  acid,  68. 


O. 

Organic  chemistry,  97. 
Oxidation  and  reduction,  73. 
Oxygen,  15. 

Oxygen  acids  of  chlorine,  43. 
Oxygen  dissolved  in  water,  167. 
Oxygen,  weight  of  a  liter  of  (quan- 
titative), 17. 
Ozone,  22. 


Permanganate  solution,  standard- 
ization of  (quantitative),  169. 
Phosphate  fertilizers,  130. 
Phosphorus,  104. 
Phosphorus,  acids  of,  107. 
Platinum,  179. 
Potassium,  122. 

R. 

Reduction,  73. 

Reduction  by  hydrogen  (quanti- 
tative), 8. 

S. 

Salts,  53. 

Salts,  acid,  basic  and  neutral,  54. 

Silicon,  98. 

Silver,  148. 

Sodium,  1 1 8. 

Sodium  chloride,  solubility  of 
(quantitative),  25. 

Solubility  in  water,  24. 

Specific  gravity  of  liquids  (carbon 
disulphide),  94. 

Specific  gravity  of  solids,  32. 

Specific  gravity  of  solids  in  small 
pieces  (sand),  101. 

Standardization  of  an  alkali  solu- 
tion (quantitative),  56. 

Strontium,  133. 


INDEX 


205 


Sulphur,  77. 
Sulphuric  acid,  82. 
Sulphurous  acid,  81. 

T. 

Tables,  193. 
Thiosulphuric  acid,  85. 
Tin,  152. 

W. 

Water,  23. 

Water,  determination  of  the  hard- 
ness of  (quantitative),  28. 


Water,  examination  of,  26. 

Water  of  crystallization  in  alum 
(quantitative),  159. 

Water  of  crystallization  in  barium 
chloride  (quantitative),  135. 

Water  of  crystallization  in  gyp- 
sum (quantitative),  31. 

Water,  oxygen  dissolved  in, 
167. 


Zinc,  139. 


A  SELECTED  LIST  OF  BOOKS  ON 

CHEMISTRY      AND       CHEMICAL 
TECHNOLOGY 

PvUithed  by 

D.    VAN     NOSTRAND    COMPANY 

25    Park    Place  New    York 


American  Institute  of  Chemical  Engineers.  Transactions. 
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BAILEY,  R.  0.  The  Brewer's  Analyst.  Illustrated.  8vo. 
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BEADLE,  C.  Chapters  on  Papermaking.  Illustrated. 
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BEAUMONT,  R.  Color  in  Woven  Design.  A  treatise  on 
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BECHHOLD,  DR.  Colloids  in  Biology  and  Medicine. 
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BEEKMAN,  J.  M.    Principles  of  Chemical  Calculations. 

In  Press. 


2  D.    VAN   NO  STRAND    COMPANY'S 

BENNETT,  HUGH  G.  The  Manufacture  of  Leather, 
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BEENTHSEN,  A.  A  Text-book  of  Organic  Chemistry. 
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BEVERIDGE,  JAMES.  Papermaker's  Pocketbook.  Spe- 
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BIRCHMORE,  W.  H.  The  Interpretation  of  Gas  Analyses. 
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BL¥TH,  A.  W.  Foods:  Their  Composition  and  Analysis. 
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Poisons :  Their  Effects  and  Detection.  A  manual  for 

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ogy. Fourth  Edition,  revised,  enlarged  and  rewritten. 
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B6CKMANN,  F.  Celluloid ;  Its  Raw  Material,  Manufac- 
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120  pp.  net,  $2.50 

BOOTH,  WILLIAM  H.  Water  Softening  and  Treatment. 
91  illustrations.  8vo.  cloth.  310  pp.  net,  $2.50 


LIST    OF    CHEMICAL   BOOKS 


BOURCART,    E.      Insecticides,    Fungicides,    and    Weed 

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BOURRY,  EMILE.  A  Treatise  on  Ceramic  Industries. 
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BRISLEE,  F.  J.  An  Introduction  to  the  Study  of  Fuel. 
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BRTJCE,  EDWIN  M.  Detection  of  the  Common  Food 
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BUSKETT,  E.  W.  Fire  Assaying.  A  practical  treatise  on 
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BYERS,  HORACE  G.,  and  KNIGHT,  HENRY  G.  Notes 
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CAVEN,  R.  M.,  and  LANDER,  G.  D.  Systematic  Inor- 
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CHRISTIE,  W.  W.  Boiler-waters,  Scale,  Corrosion,  Foam- 
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-Water,  Its  Purification  and  Use  in  the  Industries. 
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CHURCH'S  Laboratory  Guide.  A  manual  of  practical 
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4  D.    VAN   NOSTRAND    COMPANY'S 

CORNWALL,    H.    B.       Manual   of   Blow-pipe   Analysis. 

Qualitative  and  quantitative.  With  a  complete  system 
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70  illustrations.  8vo.  cloth.  310  pp.  net,  $2.50 

CROSS,  C.  F.,  BEVAN,  E.  J,  and  SINDALL,  R.  W. 
Wood  Pulp  and  Its  Uses.  With  the  collaboration  of 
W.  N.  Bacon.  30  illustrations.  I2mo.  cloth.  281 
pp.  (Van  Nostrand's  Westminster  Series.)  net,  $2.00 

d'ALBE,  E.  E.  F.  Contemporary  Chemistry.  A  survey 
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DANBY,  ARTHUR.  Natural  Rock  Asphalts  and  Bitu- 
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DEERR,  N.  Cane  Sugar.  280  illustrations.  8vo.  cloth. 
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DUMESNY,  P.,  and  NOYER,  J.  Wood  Products,  Dis- 
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DUNSTAN,  A.  E.,  and  THOLE,  F.  B.  A  Text-book  of 
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DYSON,  S.  S.,  and  CLARKSON,  S.  S.  Chemical  Works, 
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ELIOT,  C.  W.,  and  STORER,  F.  H.  A  Compendious  Man- 
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ENNIS,  WILLIAM  D.  Linseed  Oil  and  Other  Seed  Oils 
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FISCHER,  E.  Introduction  to  the  Preparation  of  Or- 
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FOYE,  J.  C.  Chemical  Problems.  Fourth  Edition,  revised 
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FRITSCH,  J.  The  Manufacture  of  Chemical  Manures. 
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GROSSMANN,  J.  Ammonia  and  Its  Compounds.  Illus- 
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HALE,  WILLIAM  J.  Calculations  in  General  Chemistry. 
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HALL,  CLARE  H.  Chemistry  of  Paints  and  Paint  Ve- 
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HILDITCH,  T.  P.  A  Concise  History  of  Chemistry. 
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HOPKINS,  N.  M.  Experimental  Electrochemistry :  Theo- 
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HOULLEVIGTTE,  L.  The  Evolution  of  the  Sciences. 
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HUDSON,  0.  F.  Iron  and  Steel.  An  introductory  text- 
book for  engineers  and  metallurgists.  With  a  section 
on  Corrosion  by  Guy  D.  Bengough.  47  illus.  8vo. 
cloth.  184  pp.  (Outlines  of  Industrial  Chemistry.) 

net,  $2.00 

HURST,  GEO.  H.  Lubricating  Oils,  Fats  and  Greases. 
Their  origin,  preparation,  properties,  uses,  and  analy- 
sis. Third  Edition,  revised  and  enlarged,  by  Henry 
Leask.  74  illus.  8vo.  cloth.  405  p.  net,  $4.00 

HYDE,  FREDERIC  S.  Solvents,  Oils,  Gums,  Waxes  and 
Allied  Substances.  554x8^.  about  200  pp.  In  Press. 

INGLE,  HERBERT.  Manual  of  Agricultural  Chemistry. 
Illustrated.  8vo.  cloth.  388  pp.  net,  $3.00 

JOHNSTON,  J.  F.  W.  Elements  of  Agricultural  Chem- 
istry. Revised  and  lewritten  by  Charles  A.  Cameron 
and  C.  M.  Aikman.  Nineteenth  Edition.  Illustrated. 
I2mo.  cloth.  502  pp.  $2.60 

JONES,  HARRY  C.  A  New  Era  in  Chemistry.  Some  of 
the  more  important  developments  in  general  chemis- 
try during  the  last  quarter  of  a  century.  Illustrated. 
i2mo.  cloth.  336  pp.  net,  $2.00 

KEMBLE,  W.  F.,  and  UNDERBILL,  C.  R.  The  Periodic 
Law  and  the  Hydrogen  Spectrum.  Illustrated.  8vo. 
paper.  16  pp.  net,  $0.50 

KERSHAW,  J.  B.  C.  Fuel,  Water,  and  Gas  Analysis,  for 
Steam  Users.  50  ill.  8vo.  cloth.  178  pp.  net,  $2.50 

KNOX,  JOSEPH.  Physico-chemical  Calculations.  I2mo. 
cloth.  196  pp.  net,  $1.00 


LIST    OF    CHEMICAL   BOOKS 


KOLLER,  T.  Cosmetics.  A  kandbook  of  the  manufac- 
ture, employment  and  testing  of  all  cosmetic  materials 
and  cosmetic  specialties.  Translated  from  the  German 
by  Charles  Salter.  8vo.  cloth.  262  pp.  net,  $2.50 

KREMANN,  R.  The  Application  of  Physico-chemical 
Theory  to  Technical  Processes  and  Manufacturing 
Methods.  Authorized  translation  by  Harold  E.  Potts, 
M.Sc.  35  diagrams.  8vo.  cloth.  215  pp.  In  Press. 

KRETSCHMAR,  KARL.  Yarn  and  Warp  Sizing  in  All 
Its  Branches.  Translated  from  the  German  by  C. 
Salter.  122  illus.  8vo.  cloth.  192  pp.  net,  $4.00 

LAMBORN.  L.  L.  Modern  Soaps,  Candles  and  Glycerin. 
224  illustrations.  8vo.  cloth.  700  pp.  net,  $7.50 

Cotton  Seed  Products.    79  illus.  8vo.  cloth.  253  pp. 

net,  $3.00 

LASSAR-COHN.  Introduction  to  Modern  Scientific 
Chemistry.  In  the  form  of  popular  lectures  suited  for 
University  Extension  students  and  general  readers. 
Translated  from  the  Second  German  Edition  by  M.  M, 
Pattison  Muir.  Illus.  I2mo.  cloth.  356  pp.  $2.00 

LETTS,  E.  A.  Some  Fundamental  Problems  in  Chemis- 
try :  Old  and  New.  44  illustrations.  8vo.  cloth.  236 
pp.  In  Press. 

LUNGE,  GEORGE.  Technical  Methods  of  Chemical 
Analysis.  Translated  from  the  Second  German  Edition 
by  Charles  A.  Keane,  with  the  collaboration  of  eminent 
experts.  To  be  complete  in  three  volumes. 
Vol.  I.  (in  two  parts).  201  illustrations.  Svo.  cloth. 
1024  pp.  net,  $1500 

Vol.  II.   (in  two  parts).     Illus.     6^x9.     1294 pp. 

net,  $18.00 
Vol.  III.  in  active  preparation. 

Technical  Chemists'  Handbook.     Tables  and  meth- 


8  D.    VAN   NOSTRAND    COMPANY'S 

ods  of  analysis  for  manufacturers  of  inorganic  chemi- 
cal products.  Illus.  I2mo.  leather.  276pp.  net,  $3.50 

Coal,  Tar  and  Ammonia.  Fourth  and  Enlarged  Edi- 
tion. In  two  volumes,  not  sold  separately.  305  illus- 
trations. 8vo.  cloth.  1210  pp.  net,  $15.00 

The   Manufacture   of   Sulphuric   Acid   and   Alkali. 

A  theoretical  and  practical  treatise. 
Vol.  I.     Sulphuric  Acid.     Fourth  Edition,  enlarged, 
In  three  parts,  not  sold  separately.     543  illustrations. 
8vo.     cloth.     1665  PP-  ne*>  $18.00 

Vol.  II.  Sulphate  of  Soda,  Hydrochloric  Acid,  Leblanc 
Soda.  Third  Edition,  much  enlarged.  In  two  parts, 
not  sold  separately.  335  illustrations.  8vo.  cloth. 
1044  PP-  net,  $15.00 

Vol.  III.  Ammonia  Soda.  Various  Processes  of  Al- 
kali-making, and  the  Chlorine  Industry.  181  illus- 
trations. 8vo.  cloth.  784  pp.  net,  $10.00 
Vol.  IV.  Electrolytical  Methods.  In  Press. 

McINTOSH,  JOHN  G.  The  Manufacture  of  Varnish  and 
Kindred  Industries.  Illus.  8vo.  cloth.  In  3  volumes. 
Vol.  I.  Oil  Crushing,  Refining  and  Boiling;  Manu- 
facture of  Linoleum  ;  Printing  and  Lithographic  Inks  ; 
India  Rubber  Substitutes.  29  illus.  160  pp.  net,  $3.50 
Vol.  II.  Varnish  Materials  and  Oil  Varnish  Making. 
66  illus.  216  pp.  net,  $4.00 

Vol.  III.  Spirit  Varnishes  and  Varnish  Materials. 
64  illus.  492  pp.  net,  $4.50 

MARTIN,  G.  Triumphs  and  Wonders  of  Modern  Chem- 
istry. A  popular  treatise  on  modern  chemistry  and 
its  marvels  written  in  non-technical  language.  76  il- 
lustrations. i2mo.  cloth.  358  pp.  net,  $2.00 

MELICK,  CHARLES  W.  Dairy  Laboratory  Guide.  52 
illustrations.  I2mo.  cloth.  135  pp.  net,  $1,25 


LIST   OF   CHEMICAL   BOOKS 


MERCK,  £.    Chemical  Reagents :  Their  Purity  and  Tests. 

8vo.    cloth.    250  pp.  New  Edition  in  Press. 

MITCHELL,  C.  A.  Mineral  and  Aerated  Waters,  in 
illustrations.  8vo.  cloth.  244  pp.  net,  $3.00 

MITCHELL,  C.  A.,  and  PRIDEAUX,  R.  M.  Fibres  Used 
in  Textile  and  Allied  Industries.  66  illustrations. 
8vo.  cloth.  208  pp.  net,  $3.00 

MUNBY,  A.  E.  Introduction  to  the  Chemistry  and 
Physics  of  Building  Materials.  Illus.  8vo.  cloth.  365 
pp.  (Van  Nostrand's  Westminster  Series.)  net,  $2.00 

MURRAY,  J.  A.  Soils  and  Manures.  33  illustrations. 
8vo.  cloth.  367  pp.  (Van  Nostrand's  Westminster 
Series.)  net,  $2.00 

NAOTET,  A.  Legal  Chemistry.  A  guide  to  the  detec- 
tion of  poisons  as  applied  to  chemical  jurisprudence. 
Translated,  with  additions,  from  the  French,  by  J.  P. 
Battershall.  Second  Edition,  revised  with  additions. 
I2mo.  cloth.  190  pp.  $2.00 

NEAVE,  G.  B.,  and  HEILBRON,  I.  M.  The  Identifica- 
tion of  Organic  Compounds.  i2mo.  cloth,  in  pp. 

net,  $1.25 

NORTH,  H.  B.  Laboratory  Notes  of  Experiments  in 
General  Chemistry.  Illus.  I2rno.  cloth,  net,  $1.00 

OLSEN,  J.  C.  A  Textbook  of  Quantitative  ChemicrJ 
Analysis  by  Gravimetric  and  Gasoinetrie  Methods. 

Including  74  laboratory  exercises  giving  the  analysis 
of  pure  salts,  alloys,  minerals  and  technical  products. 
Fourth  Edition,  revised  and  enlarged.  74  illustrations. 
8vo.  cloth.,  576  pp.  net,  $4.00 

PAKES,  W.  C.  G.,  and  NANKIVELL,  A.  T.    The  Science 

of  Hygiene.  A  text-book  of  laboratory  practice.  80 
illustrations.  I2mo.  cloth.  175  pp.  net,  $1.75 


io         D.    VAN   NOSTRAND    COMPANY'S 

PAEEY,  EENEST  J.  The  Chemistry  of  Essential  Oils 
and  Artificial  Perfumes.  Second  Edition,  thoroughly 
revised  and  greatly  enlarged.  Illustrated.  8vo.  cloth. 
554  PP-  net,  $5.00 

Food  and  Drugs.  In  2  volumes.  Illus.  8vo.  cloth. 

Vol.  I.  The  Analysis  of  Food  and  Drugs  (Chemical 
and  Microscopical).  59  illus.  724  pp.  net,  $7.50 

Vol.  II.  The  Sale  of  Food  and  Drugs  Acts,  1873- 
1907.  184  pp.  net,  $3.00 

PAETINGTON,  JAMES  R.  A  Text-book  of  Thermo- 
dynamics (with  special  reference  to  Chemistry).  91 
diagrams.  8vo.  cloth.  550  pp.  In  Press. 

—  Higher   Mathematics   for    Chemical    Students.      44 
diagrams.     i2mo.     cloth.     272  pp.  net,  $2.00 

PEEKIN,  F.  M.  Practical  Methods  of  Inorganic  Chem- 
istry. Illustrated.  i2mo.  cloth.  152  pp.  net,  $1.00 

PHILLIPS,  J.  Engineering  Chemistry.  A  pra-ctical 
treatise.  Comprising  methods  of  analysis  and  valua- 
tion of  the  principal  materials  used  in  engineering 
works.  Third  Edition,  revised  and  enlarged.  Illus- 
trated. i2mo.  cloth.  422  pp.  net,  $4.50 

PLATTNEE'S  Manual  of  Qualitative  and  Quantitative 
Analysis  with  the  Blowpipe.  Eighth  Edition,  revised. 
Translated  by  Henry  B.  Cornwall,  assisted  by  John 
H.  Caswell,  from  the  Sixth  German  Edition,  by  Fried- 
rich  Kolbeck.  87  ill.  8vo.  cloth.  463  pp.  net,  $4.00 

POLLEYN,  F.  Dressings  and  Finishings  for  Textile 
Fabrics  and  Their  Application.  Translated  from  the 
Third  German  Edition  by  Chas.  Salter.  60  illustra- 
tions. 8vo.  cloth.  279  pp.  net,  $3.00 

POPE,  F.  G.  Modern  Research  in  Organic  Chemistry. 
261  diagrams.  i2mo.  cloth.  336  pp.  net,  $2.25 


LIST    OF    CHEMICAL    BOOKS  n 

POTTS,  HAROLD  E.  Chemistry  of  the  Rubber  Industry. 
8vo.  cloth.  163  pp.  (Outlines  of  Industrial  Chem- 
istry.) net,  $2.00 

PRESCOTT,  A.  B.  Organic  Analysis.  A  manual  of  the 
descriptive  and  analytical  chemistry  of  certain  carbon 
compounds  in  common  use.  Sixth  Edition.  Illus- 
trated. 8vo.  cloth.  533  pp.  $5.00 

PRESCOTT,  A.  B.,  and  JOHNSON,  0.  C.  dualitative 
Chemical  Analysis.  Sixth  Edition,  revised  and  en- 
larged. 8vo.  cloth.  439  pp.  net,  $3.50 

PRESCOTT,  A.  B.,  and  SULLIVAN,  E,  C.  First  Book  in 
Qualitative  Chemistry.  For  studies  of  water  solution 
and  mass  action.  Eleventh  Edition,  entirely  rewritten. 
I2mo.  cloth.  150  pp.  net,  $1.50 

PRIDEAUX,  E.  B.  R.  Problems  in  Physical  Chemistry 
with  Practical  Applications.  13  diagrams.  8vo.  cloth. 
323  pp.  net,  $2.00 

PROST,  E.  Manual  of  Chemical  Analysis.  As  applied 
to  the  assay  of  fuels,  ores,  metals,  alloys,  salts,  and 
other  mineral  products.  Translated  from  the  original 
by  J.  C.  Smith.  Illus.  8vo.  cloth.  300  pp.  net,  $4.50 

PYNCHON,  T.  R.  Introduction  to  Chemical  Physics. 
Third  Edition,  revised  and  enlarged.  269  illustrations. 
8vo.  cloth.  575  pp.  $3.00 

RICHARDS,  W.  A.,  and  NORTH,  H.  B.     A  Manual  of 
Cement  Testing.      For  the  use  of  engineers  and  chem- 
ists  in   colleges   and    in   the   field.      56   illustrations. 
i2mo.     cloth.     147  pp.  net,  $1.50 

ROGERS,  ALLEN.  A  Laboratory  Guide  of  Industrial 
Chemistry.  Illustrated.  8vo.  cloth.  170  pp.  net,  $1.50 

ROGERS,  ALLEN,  and  AUBERT,  ALFRED  B.  (Editors.) 
Industrial  Chemistry.  A  manual  for  the  student  and 
manufacturer.  Written  by  a  staff  of  eminent  special- 
ists. 340  illus.  8vo.  cloth.  872  pp.  net,  $5.00 


12         D.    VAN    NOSTRAND    COMPANY'S 

ROHLAND,  PAUL.  The  Colloidal  and  Crystalloidal  State 
of  Matter.  Translated  by  W.  J.  Britland  and  H.  E. 
Potts.  I2mo.  cloth.  54  pp.  net,  $1.25 

ROTH,  W.  A.  Exercises  in  Physical  Chemistry.  Author- 
ized translation  by  A.  T.  Cameron.  49  illustrations. 
8vo.  cloth.  208  pp.  net,  $2.00 

SCHERER,  R.  Casein:  Its  Preparation  and  Technical 
Utilization.  Translated  from  the  German  by  Charles 
Salter.  Second  Edition,  revised  and  enlarged.  Il- 
lustrated. 8vo.  cloth.  196  pp.  net,  $3.00 

SCHIDROWITZ,  P.  Rubber.  Its  Production  and  Indus- 
trial Uses.  Plates,  83  illus.  8vo.  cloth.  320  pp. 

net,  $5.00 

SCHWEIZER,  V.  Distillation  of  Resins,  Resinate  Lakes 
and  Pigments.  Illustrated.  8vo.  cloth,  i83pp.net,  $3.50 

SCOTT,  W.  W.  Qualitative  Chemical  Analysis.  A  labo- 
ratory manual.  Second  Edition,  thoroughly  revised. 
Illus.  8vo.  cloth.  1 80  pp.  net,  $1.50 

SEARLE,  ALFRED  B.  Modern  Brickmaking.  260  illus- 
trations. 8vo.  cloth.  449  pp.  net,  $5.00 

SEIDELL,  A.  Solubilities  of  Inorganic  and  Organic  Sub- 
stances. A  handbook  of  the  most  reliable  quantitative 
solubility  determinations.  Second  Printing,  corrected. 
8vo.  cloth.  367  pp.  net,  $3.00 

SENTER,  G.  Outlines  of  Physical  Chemistry.  Second 
Edition,  revised.  Illus.  I2mo.  cloth.  401  pp.  $1.75 

A  Text-book  of  Inorganic  Chemistry.  90  illustra- 
tions. i2mo.  cloth.  595  pp.  net,  $1.75 

SEXTON,  A.  H.  Fuel  and  Refractory  Materials.  Second 
Ed.,  revised.  104  illus.  T2mo.  cloth.  374  pp.  net,  $2.00 

Chemistry  of  the  Materials  of  Engineering.  Illus. 

12010.  cloth.  344  pp.  net,  $2.50 


LIST    OF    CHEMICAL    BOOKS  13 

SIMMONS,  W.  H.,  and  MITCHELL,  C.  A.  Edible  Fats 
and  Oils.  Their  composition,  manufacture  and  analy- 
sis. Illustrated.  8vo.  cloth.  164  pp.  net,  $3.00 

SINDALL,  R.  W.  The  Manufacture  of  Paper.  58  illus. 
8vo.  cloth.  285  pp  .  (Van  Nostrand's  Westminster 
Series.)  net,  $2.00 

SINDALL,  R.  W.,  and  BACON,  W.  N.  The  Testing  of 
Wood  Pulp.  A  practical  handbook  for  the  pulp  and 
paper  trades.  Illus.  8vo.  cloth.  150  pp.  net,  $2.50 

SMITH,  W.  The  Chemistry  of  Hat  Manufacturing. 
Revised  and  edited  by  Albert  Shonk.  Illustrated. 
I2mo.  cloth.  132  pp.  net,  $3.00 

SOUTHCOMBE,  J.  E.     Chemistry  of  the  Oil  Industries. 

Illus.     8vo.     cloth.     209  pp.      (Outlines  of  Industrial 

Chemistry.)  net,  $3.00 

SPEYERS,  C.  L.    Text-book  of  Physical  Chemistry.      20 

illustrations.    8vo.    cloth.    230  pp.  net,  $2.25 

STEVENS,  H.  P.  Paper  Mill  Chemist.  67  illustrations. 
82  tables.  i6mo.  cloth.  280  pp.  net,  $2.50 

SUDBOROTJGH,  J.  J.,  and  JAMES,  J.  C.  Practical  Or- 
ganic Chemistry.  92  illustrations.  I2mo.  cloth. 
394  pp.  net,  $2.00 

TERRf,    H.    L.      India   Rubber   and   Its   Manufacture. 

18  illustratic-s.     8vo.     cloth.     303  pp.      (Van   Nos- 
trand's  Westminster  Series.)  net,  $2.00 

TITHERLEY,  A.  W.  Laboratory  Course  of  Organic 
Chemistry,  Including  Qualitative  Organic  Analysis. 
Illustrated.  8vo.  cloth.  235  pp.  net,  $2.00 

TOCH,  M.  Chemistry  and  Technology  of  Mixed  Paints. 
62  photo-micrographs  and  engravings.  8vo.  cloth. 
166  pp.  net,  $3.00 


14         D.    VAN   NOSTRAND    COMPANY'S 

TOCH,  M,    Materials  for  Permanent  Painting.   A  manual 

for  manufacturers,  art  dealers,  artists,  and  collectors. 
With  full-page  plates.  Illustrated.  I2mo.  cloth. 
208  pp.  net,  $2.00 

TUCKER,  J.  H.  A  Manual  of  Sugar  Analysis.  Sixth 
Edition.  43  illustrations.  8vo.  cloth.  353  pp.  $3.50 

VAN  NOSTRAND'S  Chemical  Annual,  Based  on  Bieder- 
mann's  "Chemiker  Kalender."  Edited  by  J.  C.  Olsen, 
with  the  co-operation  of  eminent  chemists.  Third 
Issue,  1913.  I2mo.  cloth.  net,  $2.50 

VINCENT,  C.  Ammonia  and  Its  Compounds.  Their 
manufacture  and  uses.  Translated  from  the  French 
by  M.  J.  Salter.  32  ill.  8vo.  cloth.  113  pp.  net,  $2.00 

VON  GEORGIEVICS,  G.  Chemical  Technology  of  Textile 
Fibres.  Translated  from  the  German  by  Charles 
Salter.  47  illustrations.  8vo.  cloth.  320  pp.  net,  $4.50 

Chemistry  of  Dyestuffs.  Translated  from  the  Sec- 
ond German  Edition  by  Charles  Salter.  8vo.  cloth. 
412  pp.  net,  $4.50 

WADMORE,  J.  M.  Elementary  Chemical  Theory.  Illus. 
I2mo.  cloth.  286  pp.  net,  $1.50 

WANKLYN,  J.  A.  Milk  Analysis.  A  practical  treatise 
on  the  examination  of  milk  and  its  derivatives,  cream, 
butter  and  cheese.  Illus.  I2mo.  cloth.  73  pp.  $1.00 

Water  Analysis.  A  practical  treatise  on  the  exami- 
nation of  potable  water.  Eleventh  Edition,  revised,  by 
W.  J.  Cooper.  Illus.  I2mo.  cloth.  213  pp.  $2.00 

WILSON,  F.  J.,  and  HEILBRON,  I.  M.  Chemical  Theory 
and  Calculations.  An  elementary  text-book.  Illus.,  3 
folding  plates.  I2mo.  cloth.  145  pp.  net,  $1.00 

WINKLER,  C.,  and  LUNGE,  G.  Handbook  of  Technical 
Gas  Analysis.  Second  English  Edition.  Illustrated. 
8vo.  cloth.  190  pp.  $4.00 


LIST    OF    CHEMICAL   BOOKS  15 

WORD  EN,  E.  C.  The  Nitrocellulose  Industry.  A  com- 
pendium of  the  history,  chemistry,  manufacture,  com- 
mercial application,  and  analysis  of  nitrates,  acetates, 
and  xanthates  of  cellulose  as  applied  to  the  peaceful 
arts.  With  a  chapter  on  gun  cotton,  smokeless  pow- 
der and  explosive  cellulose  nitrates.  Illustrated. 
8vo.  cloth.  Two  volumes.  1239  pp.  net,  $10.00 

Cellulose  Acetate.  A  monograph  of  the  history, 

chemistry,  manufacture,  technical  applications  and 
analysis  of  the  non-explosive  esters  of  cellulose  and 
starch.  Illus,  I2mo.  cloth.  In  Press. 


Any  book  in  this  list  sent  postpaid  anywhere  in  the 
world  on  receipt  of  price. 

D.    VAN    NOSTRAND    COMPANY 

Publishers  and  Booksellers 
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THIRD  THOUSAND 
868  Pages.  6^x9^  340  Illustrations.  Cloth.  NET  $5.00 

INDUSTRIAL  CHEMISTRY 

A  MANUAL  FOR  THE  STUDENT  AND  MANUFACTURER 

EDITED  BY 

ALLEN    ROGERS 

/•  than,  »f  Industrial  Cktmistry.   fntt  lutlituti.   Brooklyn.   N.  Y. 
AND 

ALFRED  B.  AUBERT 

Fonoirtr  Pnftuor  of  Chtmistry.   Vttivmity  of  Maine 

WRITTEN  BY  THESE  EXPERTS 

JEROME  ALEXANDER  JAMES  GILL1NDER  A.  H.  SABIN 

K.  H.  CLAUSSEN  W.   M.  GROSVENOR  O.  L.  SHINN 

?:l  DDO°DEOREFLINGER        ft.  I.  S5&ER  S:  k  !TELNLC£ELL 


W.   F.  FARAGHER  G.  F.  LULL  G    W    THOMPSON 

I.  C.  W.   FRAZER  I    M.  MATTHEWS  MAXMILIAN   TOCH 

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W.   K.  GANONG  OSKAR  NAGEL  O.  W    WILLCOX 

A.  H.  GILL  L.  A.  OLNEY  JOHN  H.  YOCUM 

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OPINIONS  OF  THE  PRESS 

This  book  is  "well  printed,  well  bound  and  shows  the 
results  of  careful  proof-reading.  To  get,  out  such  a  work  is  a 
great  undertaking  and  the  editors. are  to  be  congratulated 
on  the  degree  of  success  they  have  attained.  Technical  libra- 
ries, chemists,  engineers  and  teachers  will  find  this  volume 
a  necessary  addition  to  their;  works  on  industrial  chemistry. 
Journal  of  Industrial  &  Engineering  Chemistry,  Nov.,  1912 
The  editors  have  been  able  to  compile  a.  book  which  is 
absolutely  up  to  date  and  thoroughly  reliable  in  all  details 
and  which  is  of  the  greatest  value  to  everybody  who  is  inter 
ested  in  any  of  the  many  branches  of  industrial  chemistry 

The  American  Brewer^  Nov.,  1912 

The  Work  is  clear,  concise,  comprehensive  and.  modern, 
and  is  a  welcome  addition  to  the  literature  of  this  important 
branch  of  technical  knowledge. 

Drugs,  Oils  and  Paints,  Jan.,  1911 
A  need  for  such  a  volume  has  long  existed. 

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